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"V ' 







Elementary Chemistry 



THE ATOMIC THEORY— CHEMICAL COMBINATION 
—COMBINING VOLUMES OF GASES— MOLECU- 
LAR WEIGHTS OF GASES— FLAME AND THE 
COMBUSTION OF HYDRO -CARBONS- 
LIQUIDS— MATTER— PROPERTIES OF 
GASES— PROPERTIES OF SOLU- 
TIONS — VOLUME OF GASES 
—THE ELEMENTS 



/ By H. E. ROSCOE 



REVISED AND ENLARGED BY 

L. ELLIOTT BROOKES 



ILLUSTRATED 




CHICAGO 
FREDERICK J. DRAKE & CO. 

PUBLISHERS 
1907 






LIBRARY Of CONGRESS 
Two Copies Received 

JAN 12 1907 

yr Copyright Entry 

<A*~*./^ / jo L 

CUSS A XXC„ No, 

copy b. 



Copyright, 1906 

BY 

Frederick J. Drake & Co. 



PREFACE 



It has been the object of the compiler of this book to 
present the most important facts and principles of Chem- 
istry in a plain and intelligent manner, as suited to the 
requirements of students and others who may wish to 
acquire a knowledge of Elementary Chemistry. It con- 
tains full and explicit information on the following sub- 
jects : The Atomic Theory — Chemical Combination — 
Combining Volumes of Gases — Determination of Molecu- 
lar Weights of Gases — Flame and Combustion of Hydro- 
Carbons — Liquids — Matter — Properties of Gases — Prop- 
erties of Solutions — Volume of Gases — The Elements. 

The Author. 



ELEMENTARY CHEMISTRY. 



ATOMIC THEORY. 

According to the atomic theory, combination takes 
place between the atoms of which substances are made 
up, and it hence follows that the number of atoms 
which is contained in a given volume of any gaseous 
body must stand in a simple relation to that con- 
tained in the same volume of any other gas measured 
under equal circumstances of temperature and pres- 
sure. The simplest as well as the most probable sup- 
position respecting this question is that equal volumes 
of all the different gases, both elementary and com- 
pound, contain the same number of particles or mole- 
cules, and this theory is now generally accepted by 
physicists, who have arrived at the same conclusion 
as the chemists have reached by an independent train 
of reasoning. If we take the simplest case of volume 
combination, that of one volume (one molecule) of 
chlorine and one volume (one molecule) of hydrogen 
uniting to form two volumes (two molecules) of hy- 
drochloric acid gas, it is clear that, since each mole- 
cule of hydrochloric acid contains at least one atom 
of chlorine and one of hydrogen, there are at least 
twice as many atoms as molecules of these elements 
present. Hence, the molecule of free chlorine and of 
free hydrogen must consist of at least two atoms com- 

5 



6 ELEMENTARY CHEMISTRY 

bined together, and we shall represent the combina- 
tion as taking place between one volume (one mole- 
cule of two atoms) of chlorine and one volume (one 
molecule of two atoms) of hydrogen, forming two 
volumes (two molecules) of the compound hydro- 
chloric acid gas. Again, two volumes of steam are 
formed from two volumes of hydrogen and one vol- 
ume of oxygen, hence if there are the same number 
of molecules of steam, of hydrogen, and of oxygen in 
the same volume of each gas, it is clear that in the 
formation of w T ater from its elements, each molecule 
of oxygen must be split up into two similar parts. 
We are thus led to distinguish between the atom and 
the molecule, the latter term being applied to the 
smallest particle of an element or compound which 
can exist in the free state. That is, equal volumes 
of all gases, measured under the same physical con- 
ditions, contain equal numbers of molecules. 

An immediate consequence of this theory is of the 
utmost importance. If we weigh equal volumes of 
two gases, we are obviously weighing equal numbers 
of their molecules, and the ratio of the weights of the 
gases will also be the ratio of the weights of their 
molecules, so that we are thus enabled to determine 
the relative weights of the molecules of all gases, by 
simply finding their relative densities. Hydrogen gas 
is taken as the standard of comparison because it has 
a lower density than any other gas, and, since it has 
been shown that its molecule can be divided into at 
least two parts, the weight of its molecule is taken as 
equal to two. The molecular weight of any gas is 
therefore equal to twice its density compared with 
hydrogen ; nitrogen, for example, is about 14 times as 



ATOMIC THEORY 7 

heavy as hydrogen and hence its molecular weight is 
about 28, whilst carbon dioxide has a density of about 
22 and therefore a molecular weight of about 44. 

When the molecular weight of a gas and also its 
composition, as determined by analysis, are both 
known, it is possible to calculate what proportion of 
each of the component elements is present in the 
molecule. Water for instance contains 88.81 per cent 
of oxygen and 11.19 of hydrogen, whilst the density of 
steam is about 8.94, its molecular weight being, there- 
fore, equal to 17.88. If now we calculate how much 
oxygen and hydrogen are present in 17.88 parts of 
water, we find that this amount is made up of 15.88 
of oxygen and 2 of hydrogen. Carbon dioxide again 
has a molecular weight of 43.7, and contains 

Carbon 27.27 

Oxygen 72.73 



100.00 



Hence 43.7 parts of this gas contain 11.91 of carbon 
and 31.76 of oxygen. 

A repetition of this process for all the known com- 
pounds of some particular element enables us to as- 
certain the least amount of that element which is ever 
found in a molecule of a substance, and to this amount 
the name of atom is given. A comparison of all the 
substances containing oxygen, for example, teaches us 
that the least amount of it ever found in a molecule 
is 15.88 parts, and this is, therefore, taken as the 
atomic weight of oxygen. This having been ascer- 
tained, we are in a position to say that the molecule 
of water contains one atom of oxygen, whilst that of 



8 ELEMENTARY CHEMISTRY 

carbon dioxide contains two. All the non-metallic ele- 
ments form compounds which can exist in the state of 
gas, and hence the atomic weights of all these elements 
have been found by this method. Many of the metals 
on the other hand do not form volatile compounds, 
and the atomic weights of these have, therefore, to be 
determined by different methods. In any case it must 
be remembered that the method described above is not 
generally capable of great accuracy and only yields 
an approximate number for the atomic weight, the ex- 
act value being found by determining the equivalent 
of the element by an accurate analysis of one of its 
compounds, and then taking as the exact atomic weight 
the multiple of this number which approaches most 
closely to the approximate number obtained from the 
molecular weights. 

The ratio of the atomic weights of hydrogen and 
oxygen has been carefully determined by several in- 
vestigators with agreeing results, and hydrogen has 
been kept as the standard and unit, all other atomic 
weights being calculated from the ratio : H=15.88 : 1, 
found by recent experiments. 

It will be seen that the determination of the atomic 
weight is quite distinct from that of the molecular 
weight. This is well shown in the case of carbon; a 
comparison of the gaseous compounds of carbon shows 
that the atomic weight of this element is about 12, this 
being the least amount of it which is found in the 
molecule of one of its compounds, we are, however, 
quite ignorant of the molecular weight of carbon it- 
self, since its density in the state of gas has never been 
determined. 

The molecules of some elements contain as many as 



ATOMIC THEORY 9 

four atoms, phosphorus and arsenic, others contain 
only two atoms, this being the case with hydrogen, 
oxygen, nitrogen, chlorine and others, whilst the mole- 
cules of the gases of mercury vapor and of the vapors 
of some of the other metals consist of single atoms, 
the molecular and atomic weights being, therefore, 
identical. 

For the first time we may now employ chemical 
symbols, a kind of shorthand, by which we can con- 
veniently express the various chemical changes. To 
each element is given a symbol, usually the first letter 
of the Latin, which is generally also that of the Eng- 
lish name. Thus stands for oxygen, H for hydro- 
gen, S for sulphur, Au for gold (aurum), Ag for 
silver (argentum). These letters, however, signify 
more than that a particular substance takes part in 
the reaction. They serve also to give the quantity by 
weight in which it is present. Thus does not stand 
for any quantity, but for 15.88 parts by weight (the 
atomic weight) of oxygen, H always stands for one 
part of weight of hydrogen, and in like manner S, Au, 
and Ag stand invariably for 31.83, 195.7, and 107.12 
parts by weight of the several elements respectively. 
By placing symbols of any elements side by side, a 
combination of the elements is signified, thus : 

HCl=Hydrochloric acid. HI==Hydriodic acid. 
HgO=Mercuric oxide. HBr=Hydrobromic acid. 

If the molecule contains more than one atom of any 
element, this is indicated by placing a small number 
below the symbol of the atom of the element, thus H 2 
signifies 17.88 parts by weight of a compound (water) 
containing two atoms or 2 parts by weight of hydro- 



10 ELEMENTARY CHEMISTRY 

gen and one atom or 15.88 parts by weight of oxygen. 
In such a case as this, where the molecular weight and 
the number of atoms in the molecule are known and 
expressed in the formula, the latter is said to be a 
molecular formula, and consequently represents such 
a weight of the substance as will in the state of gas 
occupy the same volume as two parts by weight of 
hydrogen. When the molecular weight of the com- 
pound to be represented by a formula is not known, 
it is only possible to express the relative number of the 
atoms of the constituent elements which are present. 
A formula of this kind is known as an empirical for- 
mula and may be calculated for any substance of which 
the composition has been determined by analysis. 

The gas known as ethylene has the following com- 
position as determined by analysis : ' 

Carbon 85.62 

Hydrogen 14.38 



100.00 



In order to find the empirical formula of this sub- 
stance it is only necessary to divide the percentage of 
each element by the atomic weight of the element, 
which gives us the ratio of the number of atoms of 
each of the two elements, and if we express this ratio 
in the smallest possible whole numbers we have at 
once the relative numbers of atoms present in the 
molecule, without, however, having any information 
as to the absolute number. 

Per centag e Simplest 

Percentage. Atomic Weight. Ratio. 

Carbon 85.62 7.19 1 

Hydrogen 14.38 14.38 2 



ATOMIC THEORY 11 

The simplest or empirical formula of ethylene is 
therefore CH 2 . The density of this gas, however, is 
found to be equal to 13.91 and its molecular weight is 
therefore 27.82, its molecular formula being conse- 
quently C 2 H 4 . 

It is usual to represent chemical changes in the form 
of equations, the materials taking part in the change 
being placed on one side and the products formed, 
which are always equal to them in weight, being placed 
on the other. If we heat potassium chlorate, a sub- 
stance which has the empirical formula KC10 3 , it is 
decomposed into oxygen and potassium chloride and 
this decomposition is represented by the equation 

2KC10 3 =2KCl+30 2 

in which the sign + connects the two products and 
signifies together with. This equation is an expression 
of the fact, ascertained by experiment, that 243.36== 
(38.86+35.18+3X15.88) X2 parts of this salt by 
weight leave behind on heating 148.08 parts=(35.18+ 
38.86) X2 of potassium chloride and liberate 95.28 
parts of oxygen. Hence it is clear that the quantity 
of oxygen which is obtained from any other weight of 
the salt and vice versa can be found by a simple cal- 
culation when the equation representing the chemical 
change is known. 

To take a more complicated case, when we know 
that the equation representing the change which oc- 
curs when we heat potassium ferrocyanide, the empiri- 
cal formula of w T hich is K 4 C 6 N 6 Fe, with strong sul- 
phuric acid, H 2 S0 4 , and water, is the following: 

K 4 C 6 N 6 Fe+6H 2 S0 4 +6H 2 0==6CO+2K 2 S0 4 + 

3(NH 4 ) 2 S0 4 +FeS0 4 



12 ELEMENTARY CHEMISTRY 

yielding carbon monoxide gas CO, potassium sulphate 
K 2 S0 4 , ammonium sulphate (NH 4 ) 2 S0 4 , and iron sul- 
phate FeS0 4 , we can easily calculate how many grams 
of carbon monoxide gas, CO, can be obtained from any 
given weight of the ferrocyanide, K 4 C 6 N 6 Fe, inasmuch 
as analysis proves that the amounts represented by 
these formula are made up as follows: 

Carbon Monoxide. Ferrocyanide of Potassium. 

Carbon C 11.91 Potassium K 4 155.44 

Oxygen 15.88 Carbon C 6 71.46 

Nitrogen N 6 83.58 

27.79 Iron Fe 55.50 



365.98 



The foregoing equation then shows that 365.98 parts 
by weight of the ferrocyanide yield 166.74 parts by 
weight of carbon monoxide, and hence a simple pro- 
portion gives the quantity yielded by any other weight. 
The illustration is, however, not yet complete, com- 
mercial potassium ferrocyanide contains, as do many 
crystalline compounds, a certain quantity of water of 
crystallization, which is given off when the salt is 
heated, in consequence of which the crystals fall to a 
powder. But, as the equation shows, a certain quan- 
tity of water takes part in the reaction, and it is, there- 
fore, unnecessary to dry the salt previously if only 
we know how much water of crystallization it con- 
tains. Analysis has shown that the commercial salt 
has the composition K 4 C 6 N 6 Fe-f-3H 2 0, hence if we 
add 3X17.88, the weight of 3 molecules of water, to 
365.98, we obtain the number 419.62 as the weight of 



, ATOMIC THEORY 13 

the hydrated salt which must be taken in order to ob- 
tain 166.74 parts by weight of carbon monoxide. 

As, however, the quantity of a gas is almost always 
estimated by measuring its volume, and from this vol- 
ume calculating its weight, it becomes of the greatest 
importance to know how to calculate the volume of a 
gas from its weight, or vice versa. This can only be 
effected with strict accuracy by employing in each 
case the density of the gas as determined by experi- 
ment. An approximate number, which it is often use- 
ful to know, can, however, be readily obtained, since 
we know that molecular proportions of all gases oc- 
cupy equal volumes under the same physical condi- 
tions. 

Now 1 litre of hydrogen at 0°C. and 760 mm. pres- 
sure (which are generally taken as the standard tem- 
perature and pressure) weighs 0.089901 gram, and 
hence the volume occupied by 2 grams, or 1 gram- 
molecule of the gas is 2/0.089901=22.247 litres. The 
gram-molecule of every other gas therefore occupies 
approximately this same volume, or in other words 
the molecular volume of all gases is 22.247 litres at 
0°C. and 760 mm. pressure. 

The volume at the standard temperature and pres- 
sure of any weight of gas can then readily be calcu- 
lated, thus 27.79 grams of carbon monoxide occupy 
22.247 litres, and hence 166.74 grams occupy 

22.247X166.74 r , . Aon , _ cn 

—r^— litres at 0°C. and 760 mm. 

27.79 

It is now easy to calculate what volume this weight 
will occupy at any other temperature or pressure, for 
we know that all gases expand by T fs of their vol- 
ume at 0°C. when their temperature is raised 1°C. at 



14 ELEMENTARY CHEMISTRY 

constant pressure, and that their volume ^is inversely 
proportional to the pressure to which they are sub- 
jected. Hence if the temperature at which the gas was 
collected were 17°C, and if the barometer then stood 
a1 750 mm., the volume v in litres of the carbon monOx 
idc collected would be 

22.247X166.74X (273 + 17) X760 
27.79X273X750 



CHEMICAL COMBINATION. 

The composition of a chemical compound can be as- 
certained in two ways : By separating it into its com- 
ponent elements, an operation termed analysis, and 
by bringing the component elements under conditions 
favorable to combination, an operation termed synthe- 
sis. In both of these operations the scale is employed, 
the weight of the compound and of the components in 
each instance must be ascertained, except indeed in 
the case of certain gases of known specific gravity, 
when a measurement of the volume occupied by the 
gas may be substituted for a determination of its 
weight. 

It is one of the aims of analytical chemistry to ascer- 
tain with great precision the composition of all chem- 
ical substances, and this branch of inquiry is termed 
quantitative analysis, as distinguished from that which 
has only to investigate the kind of material of which 
substances are composed, and which is hence termed 
qualitative analysis. 

Combination by Weight. The first law discovered 
by the use of the scale, is that the elements combine 
with one another in a limited number of definite pro- 
portions, this number being almost invariably found 
by experiment to be a small one. When two elements 
are brought together under such conditions that they 
can combine, it is always found that one or more of a 
small number of compounds is produced, the particu- 
lar substance or substances formed depending upon 

15 



16 ELEMENTARY CHEMISTRY 

the special circumstances of the experiment. Thus 
carbon is found to be capable of uniting with oxygen 
in two different proportions, producing two distinct 
substances, carbonic 'acid gas and carbon monoxide, 
these being the only compounds of carbon with oxy- 
gen which are known. Some elements, on the other 
hand, only form one compound with each other, whilst 
others again form a large number. Each one of these 
compounds is found to have a fixed composition, con- 
taining the elements of which it is made up in a defi- 
nite proportion by weight, and this fixity of composi- 
tion is used as a characteristic of a chemical com- 
pound as opposed to a mere mechanical mixture, the 
constituents of which may be present in any variable 
proportions. In whatever way the conditions under 
which the elements are made to combine may be 
varied, it is always found that they unite in exactly 
the same ratio, unless, as sometimes happen, the 
changed conditions are favorable to the production of 
one of the small number of other compounds which 
can be formed by the same elements. Thus, for in- 
stance, the combination of silver with chlorine has 
been brought about in no less than four different ways, 
but in every case it has been found that the resulting 
compound contains 107.12 parts of silver for 35.18 of 
chlorine. The combination of chlorine with phos- 
phorus, on the other hand, takes place in two distinct 
ratios, so that when an excess of phosphorus is pres- 
ent, the resulting compound contains 10.25 parts of 
this element for 35.18 parts of chlorine, whilst if the 
latter be kept in excess, this weight of it only com- 
bines with 6.15 parts of phosphorus. These are, how- 
ever, the only two compounds of these elements which 



CHEMICAL COMBINATION 17 

are known. In like manner hydrogen combines with 
oxygen to yield water, a substance which contains 
88.81 parts of oxygen to 11.19 of hydrogen. If these 
elements are brought together in proportions differing 
from those in which they are present in water, the ex- 
cess of one element remains in the free state, thus, if 
98.81 parts of oxygen by weight be brought together 
with 11.19 parts of hydrogen under circumstances in 
which they can combine, 88.81 parts of the oxygen will 
combine with all the hydrogen to form 100 parts of 
water, whilst 10 parts of oxygen remain in the free 
state. 

It will therefore be seen that the chemical combina- 
tion of two or more elements does not result in the 
production of a series of compounds varying grad- 
ually in composition, according to the conditions of 
the experiment, but yields one or more compounds, 
each of which contains its constituents in a perfectly 
fixed and definite ratio. 

As has been said, the case frequently occurs of two 
elements uniting to form several compounds, for each 
of which the law of definite proportion holds good. 
Thus the two elements, carbon and oxygen, unite to 
form two distinct compounds, carbonic oxide gas and 
carbonic acid gas, and 100 parts of each of these bodies 
are found by analysis to contain the following weights 
of the elements: 

Carbonic Carbonic 

Oxide Gas. Acid Gas. 

Carbon 42.86 27.27 

Oxygen 57.14 72.73 

100.00 100.00 



18 ELEMENTARY CHEMISTRY 

Knowing these facts what is the relation of one ele- 
ment (say of the oxygen) in both compounds when 
the other element remains constant? In proportion 
to the carbon, the one compound contains exactly 
double the quantity of oxygen which the other con- 
tains, thus: 

Carbonic Carbonic 

Oxide Gas. Acid Gas. 

Carbon 10.0 10.0 

Oxygen 13.3 26.6 

23.3 36.6 



Analysis shows that two compounds which carbon 
forms with hydrogen, marsh gas and olefiant gas, have 
the following percentage composition : 



Carbon 

Hydrogen 


Marsh Gas. 

74.95 

25.05 


Olefiant Gas. 

85.68 
14.32 










100.00 


100.00 



How much hydrogen is combined in each compound 
with 10 parts by weight of carbon. In olefiant gas 
there are 1.67 parts by weight of hydrogen to 10 of car- 
bon, whilst marsh gas contains 3.34 parts of hydrogen 
to the same quantity of carbon, or exactly double as 
much. 

As another example we may take the compounds 
of nitrogen and oxygen, of which no less than five are 
known to exist. The percentage composition of these 
five bodies is found by experiment to be as follows : 



CHEMICAL COMBINATION 19 

Nitrogen 63.71 46.75 36.91 30.51 25.99 

Oxygen 36.29 53.25 63.09 69.49 74.01 



100.00 100.00 100.00 100.00 100.00 



If we inquire how much oxygen is contained in each 
of these five compounds, combined with a fixed weight, 
say 10 parts of nitrogen, we find that this is repre- 
sented by the numbers 5.7, 11.4, 17.1, 22.8, and 28.5. 
In other words, the relative quantities of oxygen are 
in the ratio of the simple numbers 1, 2, 3, 4, and 5. 

The above examples illustrate the relations exhib- 
ited in the combination of two or more of the elements 
to form compounds, but a careful examination of the 
quantitative composition of a whole series of chemical 
compounds leads to a further conclusion respecting 
the nature of the laws of chemical combination which 
is of the highest importance. Let us examine the com- 
position of any given series of compounds as deter- 
mined by analysis, such as the following: 

CHLORIDES. 

Hydrogen Chloride. Potassium Chloride. 

Chlorine 97.24 Chlorine 47.51 

Hydrogen 2.76 Potassium 52.49 



100.00 100.00 



Sodium Chloride. Silver Chloride. 

Chlorine 60.59 Chlorine 24.72 

Sodium 39.41 Silver 75.28 



100.00 100.00 



20 



ELEMENTARY CHEMISTRY 
BROMIDES. 



Hydrogen Bromide. Potassium Bromide. 

Bromine 98.76 Bromine 67.13 

Hydrogen 1-24 Potassium 32.8? 



100.00 



100.00 



Sodium Bromide. Silver Bromide. 

Bromine 77.62 Bromine 42.56 

Sodium 22.38 Silver 57.44 



100.00 



100.00 



. IODIDES. 

Hydrogen Iodide. Potassium Iodide. 

Iodine 99.21 Iodine 76.41 

Hydrogen 0.79 Potassium 23,59 



100.00 



100.00 



Sodium Iodide. 

Iodine 84.62 

Sodium 15.38 



100.00 



Silver Iodide. 

Iodine 54.03 

Silver 45.97 



100.01) 



Arranged in this way we do not notice any simple 
relation existing between the components of this se- 
except that the quantity of hydrogen is always 



ries, 



smaller than that of the chlorine, bromine, or iodine, 
whilst the quantity of sodium is always smaller than 
that of potassium, and this again is less than the quan- 
tity of silver. 

If, however, instead of examining a constant weight 
of the several compounds we ask how much of the 



CHEMICAL COMBINATION 



21 



one constituent in each compound combines with a 
constant weight of that constituent which is common 
to several, we shall obtain at once a clear insight into 
the law of the formation of the compound. In the se- 
ries of hydrogen compounds, for instance, let us cal- 
culate by simple proportion how much chlorine, bro- 
mine, and iodine combine with the unit weight of hy- 
drogen. We then obtain for the composition of these 
compounds : 

Hydrogen Chloride. Hydrogen Bromide. Hydrogen Iodide. 

Chlorine ..3518 Bromine ..79.36 Iodine ...125.90 
Hydrogen . 1.00 Hydrogen . 1.00 Hydrogen 1.00 



36.18 



80.36 



126.90 



Continuing our calculation, how much of the metals, 
potassium, sodium, and silver, unite with 35.18 parts 
by weight of chlorine to form chlorides, with 79.36 
parts of bromine to form bromides, and with 125.90 
parts of iodine to form iodides. The result is as fol- 
lows: 

CHLORIDES. 



Potassium Chloride. 

Chlorine ...35.18 
Potassium .38.86 



Sodium Chloride. 

Chlorine ...35.18 
Sodium ...22.88 



Silver Chloride. 

Chlorine ...35.18 
Silver ...107.12 



74.04 



Potassium Bromide. 

Bromine . . . 79.36 
Potassium .38.86 



58.06 
BROMIDES 

Sodium Bromide. 

Bromine ...79.36 
Sodium ...22.88 



142.30 



Sliver Bromide. 

Bromine ...79.36 
Silver ...107.12 



118.22 



102.24 



186.48 



ELEMENTARY CHEMISTRY 



Potassium Iodide. 

Iodine ...125.90 
Potassium .38.86 



IODIDES. 

Sodium Iodide. 

Iodine ...125.90 
Sodium ...22.88 



Silver Iodide. 

Iodine ...125.90 
Silver ...107.12 



164.76 



148.78 



233.02 



A remarkable relation becomes apparent, for it is 
clear that the same weights of the metals potassium, 
sodium, and silver, which combine with 35.18 parts of 
chlorine to form chlorides, also combine with 79.36 
parts of bromine to form the bromides, and with 125.90 
parts of iodine to form the iodides. In other words, 
if we replace the 35.18 parts by weight of chlorine in 
each of these compounds by 79.36 parts of bromine, we 
get the bromides, and if by 125.90 parts of iodine we 
obtain the iodides of the metals. Hence one and the 
same weight of metals, 38.86 of potassium, 22.88 of 
sodium, and 107.12 of silver, has the power of forming 
compounds with the precise quantities of chlorine, 
bromine, and iodine respectively, which unite with 1 
part by weight of hydrogen, to form the hydrides of 
these elements. These quantities of the elements in 
question are called equivalent quantities, because they 
are the amounts of them which will combine with the 
same weight of some other element. 



combine 
with 



35.18 of chlorine 
79.36 of bromine ( r !\ c 
125.90 of iodine 



tively. 



38.86 of potassium 
22.88 of sodium 
107.12 of silver 
1.00 of hydrogen 

Similar results are obtained from the examination 
of the compounds of all the other elements, so that a 
number may be assigned to each element which is 



CHEMICAL COMBINATION 23 

termed the combining weight or equivalent weight of 
the element. 

Taking an example from another group of chemical 
compounds we find that the well-known oxides of hy- 
drogen, lead, copper, mercury, and cadmium possess 
the following percentage composition : 

OXIDES. 

Water. Lead Oxide. Copper Oxide . 

Hydrogen .11.19 Lead 92.82 Copper ... .79.89 

Oxygen ...88.81 Oxygen ... 7.18 Oxygen ...20.11 



100.00 100.00 100.00 



Mercury Oxide. Cadmium Oxide. 

Mercury 92.63 Cadmium 87.54 

Oxygen 7.37 Oxygen 12.46 



100.00 100.00 



While the corresponding sulphides exhibit the fol- 
lowing composition : 

SULPHIDES. 

Sulphuretted Hydrogen. Lead Sulphide. Copper Sulphide. 

Hydrogen . 6.01 Lead 86.58 Copper ....66.48 

Sulphur .. .93.99 Sulphur . . .13.42 Sulphur . . .33.52 



100.00 100.00 100.00 



Mercury Sulphide. Cadmium Sulphide. 

Mercury 86.18 Cadmium 77.81 

Sulphur 13.82 Sulphur 22.19 



100.00 100.00 



24 ELEMENTARY CHEMISTRY 

If, as before, we now compare the quantity of each 
element united with one and the same weight of oxy- 
gen, taking 7.94 parts of this element, because this is 
the amount of it which combines with one part of hy- 
drogen and is, therefore, the equivalent weight of oxy- 
gen with respect to hydrogen, we get the following 
numbers : 

Water. Lead Oxide. Copper Oxide. 

Hydrogen ..1.00* Lead 102.68 Copper ...31.55 

Oxygen ....7.94 Oxygen .. 7.94 Oxygen ... 7.94 



8.94 



Mercury Oxide. 

Mercury 99.25 

Oxygen 7.94 



107.19 



110.62 


39.49 


Cadmium Oxide. 

Cadmium 

Oxygen 


55.SD 
7.94 




63.74 



And, if we investigate the sulphides, we find that one 
and the same weight of sulphur, 15.91 parts by weight, 
unites with weights of these elements to form sul- 
phides, which are identical with the amounts that com- 
bined with 7.94 parts by weight of oxygen to form 
oxides, thus : 

Sulphuretted Hydrogen. Lead Sulphide. Copper Sulphide. 

Hydrogen . 1.00 Lead 102.68 Copper ... .31.55 

Sulphur ...15.91 Sulphur ...15.91 Sulphur ...15.91 



16.91 118.59 47.46 



Mercury Sulphide. Cadmium Sulphide. 

Mercury 99.25 Cadmium 55.80 

Sulphur 15.91 Sulphur 15.91 

115.16 71.71 



CHEMICAL COMBINATION 



25 



Hence we see again that the amounts of these ele- 
ments which unite with an equivalent of oxygen also 
combine with an equivalent of sulphur, so that 



1.00 part by weight of hydrogen 
102.68 parts by weight of lead 
31.55 copper 

99.25 mercury 

55.80 " cadmium 



r 


m 


<x> 


■S-a 


1--0.-S \ 


a * 


o 





7.94 of oxygen 
15.91 of sulphur 



and these are, therefore, the equivalent weights of 
these elements. 

When one element combines with another in more 
than one proportion it is said to have more than one 
equivalent, and since the amounts of one element which 
combine with a fixed weight of a second are in a sim- 
ple ratio to one another, it follows that the several 
equivalents of an element must also stand in a simple 
ratio to one another. Iron for example forms several 
different compounds with oxygen, two of which have 
the following composition as determined by analysis: 

Ferrous Oxide. Ferric Oxide. 

Iron 77.54 69.97 

Oxygen 22.46 30.03 



100.00 



100.00 



Calculating the amount of iron combined with the 
equivalent (7.94 parts) of oxygen we find 

(1) (2) 

Iron 27.75 18.50 

Oxygen 7.94 7.94 



35.69 



26.44 



26 ELEMENTARY CHEMISTRY 

These amounts of iron are, however, in the simple 
ratio of 2:3, 27.75 being the equivalent of iron in fer- 
rous oxide and 18.50 in ferric oxide. 

It will be seen, therefore, that combination always 
takes place between certain definite and constant pro- 
portions of the elements or between multiples of these. 

The Atomic Theory becomes truly a chemical atomic 
theory inasmuch as it supposes the atoms of different 
elements not to possess the same weights, but to be 
characterized by different weights. Thus the atom of 
oxygen is 15.88 times as heavy as the atom of hydro- 
gen, and the weights of the atoms of oxygen and 
chlorine are as 15.88 to 35.18. Having made these 
assumptions we can understand why combination al- 
ways takes place between certain amounts of the ele- 
ments or between multiples of these amounts, since 
combination being supposed to take place between some 
number of atoms of each of the elements which unite, 
it follows that the amounts which combine must be 
some finite multiple of the weights of the atoms. 

It is thus clear that the atomic theory accounts for 
the formation of all compounds which are found to 
exist, but it is equally evident that it in no way de- 
cides how many compounds can be formed by any two 
or more elements. This at present can only be learned 
by experiment, but we are not without indications that 
a time approaches when this further problem will re- 
ceive a theoretical solution. 

COMBINING VOLUMES OF GASES. 

The relation existing between the volumes of gases 
when they combine together has been found to be a 
very simple one, inasmuch as the densities of all the 



COMBINING VOLUMES OF GASES 27 

elements known in the gaseous state are identical with 
their atomic weights, or, what is the same thing, the 
atoms in the gaseous state all occupy the same space. 

Thus the density and combining weight of oxygen 
are alike 16 (actually 15.88 with hydrogen as 1), or, 
oxygen is 16 times heavier than hydrogen, and density 
and combining weight of nitrogen are alike 14, or, 
nitrogen is 14 times heavier than hydrogen, the density 
of chlorine is 35.5, that of sulphur vapor 32, and so on. 
Remembering this fact, it is easy to calculate the ab- 
solute weight of a given volume, say one litre of these 
different gases, when we know that one litre of hydro- 
gen at the standard pressure and temperature weighs 
0.08936 grams. Thus 1 litre of oxygen, under the same 

circumstances, weighs 16X0.08936=1.430 grams 

1 litre of nitrogen weighs 14X0.08936=1.251 " 

chlorine " 35.5X0.08936=3.172 " 

sulphur vapor" 32X0.08936=2.860 " 

With respect to compounds, the density of a com- 
pound gas is one-half its molecular weight, or the mole- 
cule of a compound gas occupies the space of 2 atoms 
of hydrogen. 

18 
Thus the density of water-gas, or steam, H 2 0, is — 

A 

or 9, that is, it is nine times heavier than hydrogen, 

36 5 
the density of hydrochloric acid, HC1, is •—- or 18.25, 

A 

17 
that of ammonia, NH 3 , —or 8.5, that of carbonic acid, 

Ik 

C0 2 , ^ or 22. 

Hence the weights of 1 litre of these compounds, es- 
timated at 0° C and 760 mm., are as follows: 



28 ELEMENTARY CHEMISTRY 

1 litre of steam weighs 9X0.08936 gram 

ammonia " 8.5X0.08936 " 

hydrochloric acid " 18.25X0.08936 " 
carbonic acid " 22X0.08936 " 

The symbol for water, H 2 0, therefore, not only in- 
dicates that it is composed of 2 parts by weight of 
hydrogen and 16 of oxygen, but also that 2 volumes of 
hydrogen have united with 1 volume of oxygen to form 

2 volumes or one molecule of water gas. The symbol 
NH 3 denotes that 3 volumes of hydrogen and 1 volume 
of nitrogen have united to form 2 volumes, one mole- 
cule, of ammonia, whilst the symbol HC1 shows that 
2 volumes of hydrochloric acid gas contain 1 volume of 
chlorine and 1 of hydrogen. 

We have seen that 28 parts by weight of nitrogen 
unite with 32 parts of oxygen to form nitrogen di- 
oxide, the density of this compound is, however, found 
by experiment to be 15, hence its molecular weight is 
30, consisting of 14 parts by weight of nitrogen to 16 
of oxygen, or 1 volume of each constituent, and its 
formula must, therefore, be NO. 

Nitrogen and oxygen do not readily combine to- 
gether, but under certain circumstances they are found 
to do so, thus, if a series of electric sparks are passed 
through a glass vessel filled with dry air, the presence 
of red colored vapors, possessing a peculiar acrid smell, 
is soon noticed. These consist of nitrogen tri- and tet- 
roxides, formed by the union of the nitrogen and oxy- 
gen of the air. 

From the result of recent experiments the combining 
weight of oxygen has been taken as 15.88 instead of 
16. In the calculations used in this work, the older 
number 16 has been used in most cases for the sake of 
simplicity. 



CRYSTALLIZATION. 

Chemical substances, when they pass from the liquid 
or gaseous into the solid state, assume some definite 
geometric form, or are said to crystallize. Crystals are 
produced when a substance, such as nitre, is dissolved 
in water and the solution allowed gradually to evapo- 
rate, or when a body, such as sulphur, is melted and 
allowed to solidify by cooling. Crystalline bodies ex- 
hibit, in addition to their regular form, a peculiar 
power of splitting in certain directions more readily 
than in others, called cleavage, as well as in many cases 
the property of allowing the rays of light and heat to 
pass more readily in one direction than another, giving 
rise to the well-known phenomena of double refraction. 

Inorganic bodies which do not exhibit these pecu- 
liarities, or assume crystalline structure, are said to be 
amorphous, such as glass and glue. But certain highly 
complicated structures found in the vegetable and 
animal world exhibit a structure which, although it 
is non-crystalline, is not devoid of arrangement, and 
to w^hich the name organized or cellular structure has 
been given. 

As a rule, every substance possesses a definite form in 
which it always crystallizes, and by which it can be 
distinguished. When a crystal is formed from aqueous 
solution, the smallest visible particle possesses the com- 
plete form of the largest crystal, and simply increases 
in size without undergoing any change of form. 

Certain substances exhibiting a similarity in their 
chemical constitution are found to crystallize in the 
same forms, these are said to be isomorphous ; when the 
same body occurs crystallized in two different sys- 
tems, it is said to be dimorphous. 

29 



DETERMINATION OF THE MOLECU- 
LAR WEIGHTS OF GASES. 

For this purpose it is only necessary to determine 
the specific gravity of the gas, air being usually taken 
as the practical unit of comparison. A large glass 
balloon, capable of holding from 1 to 10 litres of the 
gas, is employed and is first of all freed from air as 
far as possible by the vacuum pump, and weighed. In 
weighing a body of such large volume it is essential to 
make allowance for the buoyancy of the air, since the 
body to be weighed appears to be lighter than it really 
is by an amount equal to the weight of the air which 
it displaces. This is best accomplished by suspending 
a vessel of similar size and shape to the other arm 
of the balance, by which arrangement the effect of 
buoyancy is neutralized, each of the vessels being af- 
fected in the same manner, and at the same time the 
uncertainties of calculation due to the varying tem- 
perature, pressure, and moisture of the atmosphere are 
avoided, as well as any inaccuracy due to condensation 
on the surface of the glass. As soon as the weight of 
the empty vessel has been ascertained, it is removed 
from the balance and filled with the pure dry gas, the 
temperature and pressure being carefully observed. 
The globe is then reweighed with the same precautions 
as before. One additional correction must be here no- 
ticed, as it gives rise to a considerable error especially 
when light gases are being weighed. The capacity of 
a vacuous globe of glass is found to be perceptibly less 
than that of the same globe when filled with gas at 

30 



MOLECULAR WEIGHT OF GASES 31 

the pressure of the atmosphere, and hence the globe 
displaces less air in the former condition than in the 
latter. The difference between the corrected weights 
of the globe empty and filled with the gas is equal to 
the weight W of the given volume of gas at the ob- 
served temperature and pressure. The same globe is 
then filled with dry air freed from carbonic acid gas, or 
else with pure hydrogen, and the weight of an equal 
volume of the latter A under the same conditions thus 
ascertained. The specific gravity of the gas compared 

W 

with air is then equal to—-. Since the air has been 

A 

found to be 14.39 times as heavy as hydrogen and the 
molecular weight of a gas is twice its density with re- 
spect to hydrogen, it is only necessary to multiply the 
specific gravity by 14.39X2 to obtain the molecular 
weight of the gas in question. 



FLAME AND THE COMBUSTION OF 
HYDROCARBONS. 

The flames which will be first considered are those 
formed when a jet of combustible gas or vapor, un- 
mixed with oxygen or inert gases, is allowed to burn in 
an atmosphere of air or oxygen. The simplest flames of 
this kind with which we are acquainted are those of 




Fig. 1. 



hydrogen and carbon monoxide. The flame of either of 
these gases burning from the end of a tube appears 
as an incandescent cone, which on examination proves 
to be hollow, the incandescence only taking place when 
the gas has mixed with air by diffusion. The hollow 
nature of these, and indeed of all flames, may be readily 
shown in various ways. A bent glass tube may be 
brought into the centre of the flame, when the unburnt 

32 



FLAME AND COMBUSTION 33 

gases will pass up the tube and may be ignited at the 
other end as in Fig. 1. The head of a match may be 
thrust quickly into the centre of the flame and held 
there for some time without the phosphorus catching 
fire, while the wood will be charred and may even take 
fire where it is in contact with the hot outer sheath of 
the cone. A thin platinum wire held horizontally in 
the flame is seen to glow at two points where it comes 
in contact with the outer zones in which the combus- 
tion is going on, whilst between them it remains cool. 

With gases which yield more than one product of 
combustion the phenomena become more complex. Thus, 
for example, the flames of cyanogen and sulphuretted 
hydrogen burning in air are found to consist of two 
sharply defined cones, possessing different colors. Mat- 
ters become still more complex in the case of the 
flames with which we are most familiar, namely, those 
obtained from a burning candle or from gaseous hydro- 
carbons. In these flames four distinct regions are 
usually distinguished ; a the dark central region, b the 
yellow region, c the blue region, d the faintly luminous 
region^ which are clearly shown in Fig. 2. 

The yellow region, which is, as a rule, the largest, 
and gives off by far the greatest amount of light, is 
known in common parlance as the luminous portion. 

The separation of carbon particles in a flame of 
hydrocarbon vapor burning in air cannot be regarded 
as due to the preferential combustion of the hydrogen, 
but is mainly owing to thermal changes taking place 
in the gas under the influence of the heat from the 
outer sheath of the flame. These changes take place in 
the dark inner cone, and finally bring about the forma- 
tion of dense hydrocarbons and carbon particles, which 



34 ELEMENTARY CHEMISTRY 

on reaching the hotter portion of the flame become in- 
candescent, giving the yellow luminous zone. 

The thermal changes which thus occur are undoubt- 
edly of a very complex nature, and succeed each other 
with such rapidity that the experimental determination 
of their exact nature is a matter of great difficulty. It 
has been shown by the analysis of the gases taken from 
a coal-gas flame at different heights that the unsatu- 







Fig. 2. 

rated hydrocarbons decrease very slowly in the dark 
portion of the flame, but quickly disappear in the lu- 
minous zone. The nature of the unsaturated hydrocar- 
bons, however, undergoes a considerable alteration in 
the non-luminous zone, the amount of acetylene in- 
creasing very rapidly and forming 70 per cent of the 
total unsaturated hydrocarbons when the top of the 
non-luminous zone is reached. The luminosity of such 
flames is due in the first place to the formation of 
acetylene by the action of heat, and the separation of 



FLAME AND COMBUSTION 35 

the carbon is due to the decomposition of the latter gas, 
which is highly endothermic, and therefore evolves heat 
in its decomposition, thus further increasing the tem- 
perature of the particles at the point of separation. 
While it is probable that the carbon separation may be 
in part brought about in this manner, it does not ap- 
pear likely that this is the sole cause of its production. 

The blue portion of the flame is the least in extent of 
any of the four divisions, and is probably caused by the 
combustion of hydrocarbons which have become mixed 
with a sufficient quantity of air to allow them to burn 
with a scarcely luminous flame, it is probable that the 
combustion is incomplete, and that the reaction going 
on here corresponds with that taking place in the inner 
cone of the flame of the Bunsen burner. The faintly 
luminous sheath is the region of complete combustion 
in which those substances which have been incom- 
pletely oxidized in the other portions of the flame, 
chiefly hydrogen and carbonic oxide, are finally con- 
verted into water and carbon dioxide. This may be 
regarded as corresponding with the outer cone of the 
Bunsen burner flame. 

When the hydrocarbon gas or vapor is mixed with 
other gases before burning, the reactions which take 
place within the flame are greatly modified, the effect 
produced depending upon the nature of the added gas. 
Such gases may consist of incombustible gases such as 
nitrogen and carbon dioxide, combustible gases such 
as hydrogen and carbon monoxide, supporters of com- 
bustion such as oxygen. 

The effect of the addition of either of the first two 
classes is to reduce the luminosity, and w^hen a suffi- 
cient percentage of the dilutent gas is added the flame 



3(3 ELEMENTARY CHEMISTRY 

becomes non-luminous, the necessary percentage vary- 
ing according to the gas used. 

The loss of luminosity thus brought about is due to 
the fact that the dilution of the hydrocarbon gas or 
vapor with other gases raises the temperature at which 
those thermal decompositions occur leading to the sep- 
aration of carbon, with the result that, before a suffi- 
ciently high temperature for such separation is reached, 
the products come in contact with the air in the outer 
flame, and there undergo complete combustion with a 
non-luminous flame. 

The luminosity of a hydrocarbon flame, unmixed 
with other gases, may also be destroyed if the flame be 
allowed to impinge on a cold surface, the temperature 
of the flame being thus reduced to such an extent that 
in this case also no separation of carbon takes place. 
This phenomenon may be readily shown by placing a 
platinum dish containing cold water over a small coal- 
gas flame, when its luminosity is destroyed, but grad- 
ually reappears as the water in the dish becomes hot. 

When the hydrocarbon gas or vapor is mixed with a 
gas which is a supporter of combustion, the flame pro- 
duced is profoundly modified. When a sufficiently 
large quantity of oxygen or air is mixed with such gas 
the flame becomes non-luminous, and at the same time 
its temperature is greatly increased. Such flames have 
a distinctly two-coned structure, the inner cone being 
pale blue, whilst the outer cone is a still paler blue. 
The most familiar example of this type of flame, which 
is of the highest technical importance, is seen in the 
Bunsen gas burner, now so universally used for heating 
purposes. In this burner the gas emerges from a cen- 
tral jet, Fig. 3, and, passing unburnt up the tube e e 



FLAME AND COMBUSTION 



37 



Fig. 4, aspirates air with it through the holes c d, the 
mixture then burning with a pale blue or bluish, smoke- 
less flame. If the air holes be closed, the gas burns 
with the usual luminous flame. 

Owing to the presence of oxygen in the gas, the ac- 
tions which now occur in the inner cone, under the 
influence of the heat from the outer sheath of flame, are 
quite different from those taking place in the corre- 





Fig. 3. 



Fig. 4. 



sponding portion of a pure hydrocarbon flame. They 
no longer consist simply of thermal decompositions, 
but of a direct combustion of the hydrocarbon with a 
limited supply of oxygen, forming products which are 
not completely oxidized, but which then burn in the 
outer cone : the inner cone becomes greener and smaller, 
and when the mixture becomes sufficiently explo- 
sive the flame does not pass down the tube as a whole, 
but the green inner cone detaches itself from the 



38 ELEMENTARY CHEMISTRY 

outer one and passes down the tube. If the tube be 
constricted lower down, the descending cone is arrested 
at that point, the speed of the gases being there greater, 
and it continues to burn there whilst the outer cone 
remains in its former position. 



LIQUIDS. 

It is evident the distinction between gas and vapor 
is only one of degree, for a vapor is simply a gas below 
its critical temperature. The same laws according to 
which the volumes of gases vary under change of tem- 
perature and pressure apply also to vapors, at any rate 
when they are examined at temperatures considerably 
above their points of condensation. When a vapor is 
near its point of condensation its density increases 
more quickly than the pressure, and as soon as the 
point is reached the least increase of the pressure 
brings about a condensation of the whole to a liquid. 

The essential difference between a liquid and the 
vapor from which it is produced lies in the fact that the 
liquid possesses a definite surface tension, and conse- 
quently occupies a definite volume limited by a surface, 
and is not capable, like a gas, of filling any space into 
which it may be brought. At the surface of the liquid 
evaporation occurs at all temperatures, and this con- 
tinues until the pressure of the vapor reaches a certain 
definite value which depends on the temperature. This 
value is known as .the vapor pressure of the liquid at 
that particular temperature, and is always reached 
when an excess of the liquid is present. When the 
vapor pressure is equal to the superincumbent pressure 
the liquid is said to boil, the corresponding tempera- 
ture being termed the boiling point of the liquid. The 
boiling points usually quoted refer to the normal at- 

39 



40 ELEMENTARY CHEMISTRY 

mospheric pressure of 760 mm. of mercury, and an at- 
mospheric temperature of 0° Centigrade. 

Liquids possess a notable vapor pressure below their 
boiling-points, thus water gives off vapor at all tem- 
peratures, and even slowly evaporates when in the 
solid state, for the pressure of the vapor coming from 
ice at — 10° is 0.208 mm. According to experiments 
there is a limit beyond which evaporation cannot be 
detected, thus mercury gives out a perceptible amount 
of vapor during the summer, but that none can be de- 
tected during the winter, and certain compounds which 
can be volatilized at 150°, undergo no perceptible 
evaporation when kept for years at the ordinary tem- 
perature. This maximum pressure or density of a 
vapor is not altered by the presence of other gases, or, 
in other words the quantity of a liquid which will 
evaporate into a given space is the same whether the 
space is a vacuum or is filled with another gas. The 
above conclusions can only be considered as approxi- 
mately true, as about 2 per cent more vapor ascends 
into a space filled with gas than into a vacuum, while 
at considerable but equal distances from the boiling- 
point the pressures of volatile liquids are by no means 
equal. 



MATTER. 

Matter is capable of assuming three different states 
or conditions: the solid, the liquid, and the gaseous. 
Of these, the first two have, for obvious reasons, been 
recognized from the earliest ages, as accompanying 
very different kinds of substances. It is, however, only 
within a comparatively short time that we have come 
to understand that just as there are many distinct 
kinds of solids and liquids, so there are many distinct 
kinds of gases. These may be colorless and invisible, 
but they can readily be skow r n to differ one from an- 
other. A peculiar gas, which we now know as carbonic 
acid gas, or carbon dioxide, was obtained by the action 
of dilute acids on marble, to this gas was given the 
name of fixed air, because it is fixed in the alkaline 
carbonates, which at that time were called the mild al- 
kalies, in contradistinction to the caustic alkalies. This 
invisible gas does not, like air, support the combustion 
of a taper, and, unlike air, it renders clear lime-water 
turbid, it is also much heavier than air, as can be 
shown by pouring it downwards from one vessel to 
another, by drawing it out of a vessel by means of a 
syphon, or by pouring it into a beaker glass previously 
equipoised at one end of the beam of a balance as in 
Fig. 5. That the gas has actually been poured out is 
seen either by a burning taper being extinguished 
when dipped into the beaker glass, or by adding some 
clear lime-water, which then turns milky. 

41 



42 



ELEMENTARY CHEMISTRY 



The gas which was formerly termed inflammable air, 
and obtained by the action of dilute acids on metallic 
zinc or iron, is also a peculiar and distinct substance, 
to which we now give the name of hydrogen gas. It is 




Fig-. 5. 



so much lighter than air that it may be poured up- 
wards, and takes fire when a light is brought in con- 
tact with it, burning with a pale blue flame. Soap- 
bubbles blown with hydrogen ascend in the air, and if 



MATTER 



43 



hydrogen be poured upwards in the equipoised bell- 
jar hung mouth downwards on the arm of the balance 
as in Fig. 6, the equilibrium will be disturbed, and the 
arm with the bell-jar will rise. 




Fig. 6. 



Chemists have come to the conclusion that matter is 
indestructible and that in all eases of chemical action 
in which matter disappears, the loss is apparent only, 
the solid or liquid being changed into an invisible gas, 



44 ELEMENTARY CHEMISTRY 

the weight of which is, however, exactly identical with 
that of its component parts. We only require to allow 
a candle to burn for a few minutes in a clean flask 
filled with air in order to show that the materials of 
the candle, hydrogen and carbon, unite with the oxygen 
of the air to form, in the first place, water, which is 
seen in small drops bedewing the bright sides of the 
flask, and in the second, carbon dioxide or carbonic 
acid gas, whose presence is revealed to us by lime- 
water being thereby turned milky. The fact that the 
sum of the weights of the products of combustion (wa- 
ter and carbon dioxide) is greater than the loss of 
weight sustained by the candle is clearly shown by an 
experiment made by means of the apparatus shown in 
Fig. 7, which consists of a tube equipoised on the arm 
of a balance. In the long vertical tube a taper is 
placed, the other end of the system being attached to 
a gasholder filled with water, which, on being allowed 
to run out, causes a current of air to pass through the 
tube, and thus maintains the combustion of the taper. 
The water and carbonic acid gas which are formed are 
absorbed in the bent tube, which contains caustic 
potash. After the taper has burnt for a few minutes, 
the apparatus is disconnected from the gasholder and 
allowed to vibrate freely, when it will be found to be 
appreciably heavier than it was before the taper had 
burnt, the explanation being that the excess of weight 
is due to the combination of the carbon and hydrogen 
of the wax with the oxygen of the air. 

The statement that matter is indestructible is based 
entirely upon the evidence of experiment, and many 
investigations have been carried out to test its validity. 
These have taken the form of weighing two substances, 



MATTER 45 

such as silver and iodine^ mercury and iodine, or iodine 
and sodium sulphite, as carefully as possible, and then 
allowing them to unite or react chemically and finally 
weighing the products. In the experiments the silver 
and iodine were separately dissolved and allowed to 



Fig. 7. 

react, and the resulting silver iodide was then collected 
and weighed, but in the other cases the chemical change 
was allowed to occur within sealed-up vessels, so that 
no mechanical loss could take place. The result has 
been that no definite change of weight has in any case 
been observed. The accuracy of the statement that 



46 ELEMENTARY CHEMISTRY 

matter is indestructible is therefore true within the 
limits of accurate weighing which have at present been 
attained. 

It is the aim of the chemist to examine the properties 
of all the different substances which occur in nature, so 
far as they act upon each other, or can be made to act 
so as to produce something different from the sub- 
stances themselves, to ascertain the circumstances un- 
der which such chemical changes occur, to discover the 
laws upon which they are based, and to investigate the 
relation between the properties of substances and their 
chemical composition. In thus investigating terrestrial 
matter it is found that all the various forms of matter 
with which we are surrounded, or which have been 
examined, can be divided into two great classes. 

Elementary Bodies. Elements, or simple substances, 
out of which no other two or more essentially 'differing 
substances have been obtained. 

Compound Bodies, or compounds, out of which two 
or more essentially differing substances have been ob- 
tained. 

Of these, and their compounds with each other, the 
whole mass of our globe, solid, liquid, and gaseous, is 
composed. 



MATTER 



47 



List of Elements. 




Atomic Weight. 




Atomic Weight. 




H=l. 






H=l. 


Aluminum 


Al 


26.9 


Neon 


...Ne 


19.9 


Antimony 


....Sb 


119.3 


Nickel 


....Ni 


58.3 


Argon 


A 


39.6 


Nitrogen 


N 


13.93 


Arsenic 


As 


74.4 


Osmium : 


...Os 


189.6 


Barium 


Ba 


136.4 


Oxygen 


O 


15.88 


Bismuth 


Bi 


206.9 


Palladium 


...Pd 


105.7 


Boron 


B 


10.9 


Phosphorus .. 


P 


30.77 


Bromine 


Br 


79.36 


Platinum 


....Pt 


193.3 


Cadmium 


Cd 


111.6 


Potassium 


K 


38.86 


Caesium 


Cs 


131.9 


Praseodymium 


L..Pr 


139.4 


Calcium 


Ca 


39.8 


Radium 


..Ra 


223.3 


Carbon 


C 


11.91 


Rhodium 


...Eh 


102.2 


Cerium 


Ce 


139.2 


Rubidium 


...Rb 


84.8 


Chlorine 


CI 


35.18 


Ruthenium 


...Ru 


100.9 


Chromium 


Cr 


51.7 


Samarium 


..Sm 


148.9 


Cobalt 


Co 


58.56 


Scandium 


...Sc 


43.8 


Copper 


Cu 


63.1 


Selenium 


....Se 


78.6 


Erbium 


Er 


164.8 


Silicon 


Si 


28.2 


Fluorine 


F 


18.9 


Silver 


-Ag 


107.12 


Gadolinium .... 


Gd 


155. 


Sodium 


...Na 


22.88 


Gallium 


...Ga 


69.5 


Strontium 


....Sr 


86.94 


Germanium .... 


.....Ge 


71.9 


Sulphur 


S 


31.83 


Gold 


....Au 


195.7 


Tantalum 


...Ta 


181.6 


Helium 


....He 


4. 


Tellurium 


....Te 


126.6 


Hydrogen 


H 


1.000 


Terbium 


...Tb 


158.8 


Indium 


In 


131.1 


Thallium 


....Tl 


202.6 


Iodine 


I 


125.90 


Thorium 


...Th 


230.8 


Iridium 


Ir 


191.5 


Thulium 


..Tm 


169.7 


Iron 


Fe 


55.5 


Tin 


...Sn 


118.1 


Krypton 


Kr 


81.2 


Titanium 


...Ti 


47.7 


Lanthanum 


La 


137.9 


Tungsten 


...W 


182.6 


Lead 


...Pb 


205.35 


Uranium 


U 


236.7 


Lithium 


Li 


6.98 


Vanadium 


V 


50.8 


Magnesium 


...Mg 


24.18 


Xenon 


X 


127. 


Manganese 


....Mn 


54.6 


Yttrbium 


...Yb 


171.7 


Mercury 


Hg 


198.5 


Yttrium 


...Yt 


88.3 


Molybdenum.. 


....Mo 


95.3 


Zinc 


...Zn 


64.9 


Neodymium .. 


....Nd 


142.5 


Zirconium 


....Zr 


89.9 



PROPERTIES OF GASES. 

The chemist has to deal with matter in all its various 
states, solids, liquids, and gases alike being the objects 
of his examination. The study of gases in particular 
has, as we have seen, led to most important results in 
the theoretical branch of the science, the system of 
formulae now employed being in fact founded upon ob- 
servations made upon matter in the state of gas. 

Boyle's Law. The gaseous condition of matter is 
well defined to be that in which it is capable of indefi- 
nite expansion. If a quantity of gas as small as we 
please be placed in a closed vacuous space, however 
large, the gas will distribute itself uniformly through- 
out that space. The relation between the volume and 
pressure of a gas, the temperature remaining constant, 
is expressed by the well-known law of Boyle, that the 
volume of the gas varies inversely as the pressure, 
from which it follows that the product of pressure and 
volume remains constant whatever be the pressure. 
This is expressed by the equation PV=C, thus when 
the pressure is doubled, the volume is halved, the prod- 
uct of the two remaining the same. When a number 
of gases which do not act chemically on each other are 
mixed together, the total pressure exerted is equal to 
the sum of the separate pressures which each gas would 
exert if it alone occupied the whole space, or as it may 
be otherwise expressed, the total pressure of a mix- 
ture of gases is equal to the sum of the partial pres- 
sures of its constituent gases. 

48 



PROPERTIES OF GASES 49 

Thus, if a litre of oxygen at a pressure of 0.2 of an 
atmosphere and a litre of nitrogen at a pressure of 0.8 
of an atmosphere be mixed and the volume again 
brought to 1 litre, the total pressure will be 1 atmos- 
phere, this being the sum of the partial pressures of 
the oxygen and nitrogen. 

The adoption of the pressure of a column of mer- 
cury 760 mm. high as the normal or standard pressure 
leads to the anomaly that the mass of a given volume 
of gas, under standard conditions, varies at different 
places on the earth's surface, since the pressure 
exerted by a column of 760 mm. of mercury varies with 
the latitude and the height above sea level, owing to 
the variation in the intensity of gravitation. Hence 
when the weight of a given volume is quoted the local- 
ity for which it has been determined must also be 
stated. 



PROPERTIES OF SOLUTIONS. 

A remarkable anajogy exists between the properties 
of substances in dilute solution and those of gases. This 
conclusion has been chiefly derived from the study 
of the interesting phenomena now to be described. 
When a solution of a crystalloid substance in water is 
placed in a vessel closed by a porous membrane, such 
as a piece of parchment, and the whole immersed in 
pure water, it is found that the dissolved substance 
gradually passes outwards through the film of parch- 
ment, whilst water passes inwards, until, after a suffi- 
cient time has elapsed, equilibrium is established and 
the liquid has the same composition both inside and 
outside the membrane, this process being known as 
osmosis. 

Diaphragms of other substances can, however, be 
obtained which allow the water to pass freely through 
them in the same way as the parchment, but prevent 
the outward passage of the dissolved substance, and are 
therefore said to be "semipermeable." Such a dia- 
phragm can be prepared by filling an ordinary porous 
cell with a dilute solution of potassium ferro-cyanide 
and simultaneously immersing it in one of copper sul- 
phate. These two substances gradually diffuse into the 
porous walls of the cell and produce an insoluble layer 
of copper ferrocyanide, which is found to be semi- 
permeable for solutions of many salts and other bodies. 
If now such a cell, containing a dilute solution of sugar, 
be placed in a vessel containing pure water, the latter 

50 



PROPERTIES OF SOLUTIONS 51 

is found to pass inwards through the film, whilst the 
sugar does not pass out, and consequently the level 
of the liquid within the cell rises. If, however, the cell 
be completely filled and connected with an arrangement 
for measuring the pressure it will be found that the 
latter gradually increases, but after a time becomes 
constant. The pressure thus observed is termed the 
osmotic pressure of the solution, and is found to de- 
pend only upon the nature of the dissolved substance, 
the concentration of the solution, and the tempera- 
ture. 

This osmotic pressure plays the same part in the 
theory of dilute solutions as the gaseous pressure in 
that of gases, and is found to follow the same laws. 
Thus when the strength of the solution is doubled the 
osmotic pressure becomes twice as great, when it is 
halved the pressure falls to one-half, and so on. Now 
doubling the strength of a solution is in reality halving 
the volume occupied by the unit weight of the dis- 
solved substance, so that the law that the osmotic 
pressure of a dilute solution varies directly as its con- 
centration corresponds exactly with Boyle's law of 
gases. 



VOLUME OF GASES. 

Gases are known as compressible fluids, and liquids 
as incompressible fluids. Liquids really are compres- 
sible, but only to a very slight extent. Like gases, they 
recover volume on removal of the pressure. The law 
representing the relation between the volumes of a 
gas and the pressures to which the gas is subjected is 
a very simple one. It is termed Boyle's or Mariotte's 
Law, from the names of the discoverers. It states that 
the volume occupied by any gas is inversely propor- 
tional to the pressure to which it is subjected. Thus 
the volume 1 under pressure 1 becomes the volume 2 
under the pressure %, the volume 3 under the pressure 
y s , the volume % under the pressure 2, and the volume 
% under the pressure 3, and so on. 

The instrument which serves to measure the pres- 
sure exerted by the air is termed a barometer. This 
in its simplest form consists of a straight glass tube, 
about 800 mm. (33 inches) in length, closed at one end, 
and furnished with a millimeter scale. This tube is 
filled with mercury, and the open end placed down- 
wards in a basin containing the same metal. It is then 
seen that the mercury sinks in the tube to a point 
about 760 mm. from the surface of the metal in the 
basin. It is sustained in this position by the pressure 
of the air. When this pressure increases, the height 
of the sustained column becomes greater when it 
diminishes, the level of the mercury in the tube falls. 
All gases generated at the earth's surface are subject 

52 



DIFFUSION OF GASES 53 

to this pressure, and* their volumes increase or diminish 
according to the above law, as the superincumbent 
pressure becomes less or greater. In estimating the 
volume of hydrogen which can be collected from a 
given weight of zinc and sulphuric acid, it is clear that 
we require to know not only the temperature at which 
the gas is collected, but also the atmospheric pressure 
under which it is measured, and in order to be able to 
compare the bulks of two gases, we must always com- 
pare them under like conditions of temperature and 
pressure. For this purpose we compare all the volumes 
of gases at the standard temperature of 0° C, and un- 
der the standard pressure of 760 millimetres of mer- 
cury. 

DIFFUSION OF GASES. 

Another physical property of gases is that of diffu- 
sion. Gases which, when mixed together, do not com- 
bine chemically, have the power of becoming intimately 
mixed together, even when differing in specific gravity, 
and when the heavier gas is placed at the bottom, and 
both remain at rest. This important property is called 
the diffusive power of gases. The rate at which gases 
diffuse varies greatly. Thus, a bottle filled with hydro- 
gen lost 94.5 per cent of this gas when left exposed to 
the air in the same time as that in which a bottle of car- 
bonic acid lost only 47 per cent of this gas in the same 
way. Gaseous diffusion goes on through the minute 
pores of certain solids, such as stucco, or thin plates of 
graphite; the different diffusive rates of air and hy- 
drogen may be well seen by fixing a thin piece of stucco 
on to one end of a glass tube open at the other end, 
and filling this with hydrogen, on plunging the open 



54 



ELEMENTARY CHEMISTRY 



end into water a steady rise of this liquid in the tube is 
noticed, and after some time the whole of the hydrogen 
is found to have disappeared, and the tube contains 
only pure air. Experiments made upon this subject 
have shown that the velocity of diffusion of different 
gases is inversely proportional to the square roots of 
their densities, thus 4 volumes of hydrogen will pass 
through the diaphragm in the same time that 1 volume 
of oxygen is able to do so, oxygen being approximately 
sixteen times as heavy as hydrogen. This property of 
gases has an important bearing upon the atmosphere 
of towns and dwelling-rooms, which is kept pure to a 
great extent by this diffusive power of gases. 

The following table gives the rates of diffusion of 
several gases, that of air taken to be equal to 1, com- 
pared with the inverse square roots of their densities, 
air also taken as their unit : 



Gas. 


Density 
air=l. 


1 


Velocity of 

diffusion 

air=l. 




v density 


Hydrogen 

Nitrogen 

Oxygen 

Carbon Dioxide 


0.06926 
0.97130 
1.10560 
1.52900 


3.7790 
1.0150 
0.9510 

0.8087 


3.830 
1.014 
0.949 

0.812 





THE ELEMENTS. 



ALUMINUM. 
Symbol— Al. Atomic Weight— 26.9. 

This metal occurs in large quantities combined with 
silicon and oxygen in felspar and all the older rocks, 
and also in clay, marl, slate, and in many crystalline 
minerals. Metallic aluminum is obtained by passing 
the vapor of aluminum chloride over metallic sodium. 
It has recently been manufactured on a large scale and, 
from its lightness and its bright lustre, it has been 
used for the metallic portions of optical instruments as 
well as for ornamental work. 

Clay is an aluminum silicate resulting from the dis- 
integration and decomposition of felspar by the action 
of air and water, the soluble alkali being washed away. 
Kaolin or porcelain clay is the purest form of disinte- 
grated felspar, containing no iron or other impurities. 
There are many very beautifully crystalline minerals, 
consisting of aluminum silicates combined with silicates 
of the metals of the alkalies and alkaline earths, among 
others, garnet, idocrase, and mica. 

Aluminum salts can be detected when in solution by 
giving with ammonia a white precipitate, insoluble in 
excess, but soluble in caustic soda, and by assuming a 
blue color when moistened with cobalt solution and 
heated before the blowpipe. 

55 



56 ELEMENTARY CHEMISTRY 

Aluminum and Oxygen Compounds, 

Alumina : Formula — A1 2 3 . 

Alumina is the only oxide of aluminum known. It 
occurs native in a nearly pure and crystalline state as 
corundum, ruby, sapphire, and in a less pure state as 
emery. Alumina is prepared by adding ammonia to a 
solution of alum, a white precipitate of the hydroxide 
falls down, and this on being heated yields a white 
amorphous powder of pure alumina. This substance is 
attacked with difficulty by acids, but the hydrate is 
easily soluble in acids and in the fixed caustic alkalies. 
Alumina acts as a weak base, the commonest aluminum 
salts are the alums, and their solutions have an acid 
reaction. Alumina is largely used in dyeing and calico- 
printing as a mordant, as it has the power of forming 
insoluble compounds called lakes with vegetable color- 
ing matter, and thus renders the color permanent by 
fixing it in the pores of the cloth so that it cannot be 
washed out : such colors are termed fast. 

Aluminum and Chlorine Compounds. 
Aluminum Chloride: Formula — A1 2 C1 6 . 

Aluminum chloride is a volatile white solid body, 
obtained by heating a mixture of alumina and charcoal 
in a current of chlorine gas, it is used in the manufac- 
ture of the metal. 

Aluminum, Sulphur and Oxygen Compounds. 
Aluminum Sulphate : Formula — A1 2 3S0 4 . 

Aluminum sulphate is a soluble salt prepared on a 
large scale for the use of the dyer by decomposing 
clay, by acting upon it with sulphuric acid. The solid 



ANTIMONY 57 

mixture of silica and aluminum sulphate thus obtained 
goes by the name of alum-cake. The most useful com- 
pounds of alumina are, however, the alums, a series of 
double salts, which aluminum sulphate forms with the 
alkaline sulphates. A large number of other alums 
are known, in which the isomorphous sesquioxides of 
iron, chromium, and manganese are substituted for the 
alumina in common alum. 

ANTIMONY. 

Symbol— Sb. Atomic Weight— 119.4. 

Metallic antimony occurs native, but its chief ore is 
the trisulphide. The metal is easily reduced by heating 
the sulphide with about half its weight of metallic 
iron, when ferrous sulphide and metallic antimony are 
formed. Antimony may also be reduced by mixing the 
ore with coal and heating in a reverberatory furnace. 
Antimony is a bright bluish-white colored metal, it is 
very brittle, and can be powdered in a mortar, it melts 
at 450°, and may be distilled at a white heat in an 
atmosphere of hydrogen. Antimony undergoes no al- 
teration in the air at ordinary temperatures, but rapid- 
ly oxidizes if exposed to air when melted, and, if heated 
more strongly, it takes fire and burns with a white 
flame, giving off dense white fumes of antimony tri- 
oxide. Antimony is not attacked either by dilute hy- 
drochloric or sulphuric acids. Nitric acid attacks the 
metal, converting it into white insoluble antimony 
pentoxide. Nitro-hydrochloric acid dissolves antimony 
easily. The alloys of antimony are largely used in the 
arts. Of these, type metal, an alloy of lead and anti- 
mony, is the most important, it contains 17 to 20 per 
tent of the latter metal. 



58 ELEMENTARY CHEMISTRY 

Antimony and Oxygen Compounds. 

Antimony Trioxide, Sb 2 (X and Antimony Pentoxide, 
Sb 2 5 . 

This oxide gives rise to the important series of salts 
of antimony used in medicine, it is obtained in crystal- 
line needles, which are isomorphous with the rare form 
of arsenic trioxide. Antimony trioxide has also been 
observed to crystallize in octahedra, hence these two 
oxides are said to be iso-dimorphous. The best mode 
of preparing the pure, oxide is by decomposing anti- 
mony trichloride with an alkaline carbonate, when the 
oxide is precipitated as a white powder. 

Antimony trioxide dissolves, when boiled with a solu- 
tion of cream of tartar (hydrogen potassium tartrate), 
and on concentration the solution deposits crystals of 
tartar emetic (potassium antimony tartrate), antimony 
trioxide also dissolves in hydrochloric acid, yielding a 
solution of the trichloride, which is rendered turbid by 
addition of water, owing to the formation of an in- 
soluble antimony oxychloride. 

Antimony Pentoxide: Formula — Sb 2 5 . 

Antimony Pentoxide, sometimes called Antimonic 
Acid, obtained by acting on antimony with strong ni- 
tric acid, or by decomposing the pentachloride of anti- 
mony with water. It is a light straw-colored powder, 
which loses oxygen at a red heat, and is converted 
into the intermediate oxide Sb 2 3 Sb 2 3 . Antimony 
pentoxide forms salts with the alkalies called antimoni- 
ates, from which antimonic acid, Sb 2 6 H 2 , can be 
separated as a white powder. The oxides prepared by 
the two methods above given are found to possess dif- 
ferent properties as regards their power of uniting 



ARSENIC 59 

with bases. That prepared with nitric acid yields 
monobasic salts, while that obtained from the penta- 
chloride yields dibasic salts. To the first class of salts 
the name antimoniates, and to the second that of 
metantimoniates, has been given. 

The grey intermediate tetroxide, Sb 2 4 , is obtained 
by heating the metal in the air until no further change 
occurs. 

Finely-powdered metallic antimony takes fire sponta- 
neously when thrown into chlorine gas, with formation 
of the chlorides. 

ARSENIC. 

Symbol — As. Atomic Weight — 74.4. 

Arsenic closely resembles phosphorus in its chemical 
properties and in those of its compounds, although in 
physical characters, such as specific gravity and lustre, 
it bears a greater analogy to the metals. It may be 
considered the connecting link between these two divi- 
sions of the elements, antimony and bismuth being 
closely connected with it on the one hand, and phos- 
phorus and nitrogen on the other. Arsenic is some- 
times found in the free state, but more frequently com- 
bined with iron, nickel, cobalt, and sulphur. It is also 
contained in very small quantities in many mineral 
springs. In order to separate arsenic from any of the 
metallic ores in which it occurs, the ore is roasted, or 
exposed to a current of heated air in a reverberatory 
furnace, the arsenic combines with the atmospheric 
oxygen, forming arsenic trioxide, As 2 3 , which is car- 
ried in the state of vapor from the furnace into long 
chambers or flues, in which the trioxide, commonly 



60 ELEMENTARY CHEMISTRY 

known as arsenious acid, or white arsenic, is deposited. 
Metallic arsenic may be prepared from this oxide by 
mixing it with charcoal and sodium carbonate, and 
heating in a closed crucible, the upper part of which 
is kept cool, arsenic condenses in the cool part of this 
apparatus as a solid with a brilliant greyish lustre. It 
tarnishes in the air from oxidation, it has a specific 
gravity of 5.7 to 5.9, and when heated to dull red- 
ness, it volatilizes as a colorless vapor without under- 
going fusion, and this vapor possesses a remarkable 
garlic-like smell. Arsenic when heated in the air takes 
fire, and burns with a bluish flame, forming arsenic 
trioxide As 2 O s , when thrown into chlorine, it instantly 
takes fire, forming arsenic trichloride, AsCl 3 . 

In order to purify the commercial arsenic it is sub- 
limed with the addition of a small quantity of pow- 
dered charcoal. On a small scale, arsenic may be 
purified by introducing the mixture into a glass flask, 
which is then placed in a large crucible, surrounded by 
sand, and heated t6 redness. As soon as the sublima- 
tion begins, a loosely-fitting stopper of chalk is placed 
in the neck of the flask, a second crucible is placed over 
the first, and the whole heated until the arsenic is sub- 
limed into the upper portion of the flask. In this way 
crystals of arsenic are obtained, which have a bright 
metallic lustre. 

Arsenic has a steel gray color, and is a good conduc- 
tor of electricity. If pure arsenic is quickly sublimed 
in a stream of hydrogen gas it is deposited in the 
neighborhood of the heated portion of the tube in crys- 
tals, but, at a little distance, as a black glittering mass, 
• and, still further on, as a yellow powder. Both the 
black and the yellow modifications are crystalline, Yel- 



ARSENIC 61 

low arsenic is best prepared by distilling arsenic in 
a stream of carbon dioxide, the vapors being passed 
into a U-tube, where the arsenic is condensed by com- 
ing into contact with another stream of carbon dioxide 
which has been cooled, the arsenic is then dissolved 
in carbon bisulphide, from which it is deposited by 
evaporation or by cooling to — 70°. 

This form of arsenic is extremely sensitive to light, 
changing quickly into the ordinary one, but a solution 
in carbon bisulphide may be kept for some time un- 
changed, whilst both the yellow and the black modifi- 
cations give metallic arsenic on heating. 

A reddish-brown crystalline modification, which is 
not changed by light, is stated to be deposited from 
a solution of the yellow arsenic in carbon bisulphide 
on long standing. 

It was formerly supposed that arsenic could not be 
melted, for when heated under ordinary circumstances 
to about 450° it passes at once from the solid to the 
gaseous state. Its melting point lies between those of 
antimony and silver. 

Arsenic and Oxygen Compounds. 

Arsenious Oxide: Formula — As 4 6 . 

Molecular Weight— 361.1. 

Arsenious oxide has long been known under the 
names of white arsenic and arsenious acid. The name 
of smelting-furnace-smoke is given to this substance 
because it is obtained by roasting arsenical pyrites, 
and is emitted during the process in the form of a 
white smoke which condenses to a white powder. 

Arsenious oxide is prepared on a large scale in many 
metallurgical processes by the roasting of arsenical 



62 ELEMENTARY CHEMISTRY 

ores. The vapors of the oxide which are given off are 
condensed in long passages or chambers called poison- 
chambers, formerly in towers termed poison-towers, 
in the form of crude flowers of ars£nic or poison-flow- 
er. For the preparation of white arsenic, arsenical 
pyrites is usually employed. It is obtained as a by- 
product in the roasting of cobalt ores, which are em- 
ployed in the manufacture of smalt. When kept for 
any length of time it becomes opaque, being 
changed into a porcelain-like mass. This change is 
due to the passage from the amorphous or vitreous to 
the crystalline condition, the change commences at 
the outside of the mass, and gradually penetrates into 
the interior. 

Arsenious oxide is slightly soluble in w^ater. It is 
very much more soluble in hydrochloric acid than in 
water, and it may be easily obtained from the solution 
in the form of large crystals. It also occurs in this 
form as arsenic-bloom, being found together with na- 
tive arsenic, having been formed by the oxidation of 
this substance. 

Arsenious oxide serves for the preparation of a large 
number of other arsenic compounds, especially of the 
acids of arsenic and their salts. It is also employed 
in the manufacture of arsenical pigments and is large- 
ly used in the manufacture of glass. 

Arsenic and Sulphur Compounds. 

Arsenic subsulphide, As 3 S, Arsenic disulphide, As 2 S 2 

and Arsenic trisulphide, As 2 S 3 . 

Arsenic subsulphide: Formula — As 2 S. 

This compound is obtained by adding phosphorus 
trichloride to an aqueous solution of sodium arsenate, 



ARSENIC 63 

allowing the mixture to cool, saturating with sulphur 
dioxide, and allowing the whole to stand for two or 
three days, when a dark brown precipitate of arsenic 
subsulphide separates out. 

This sulphide 2 on heating, decomposes, giving real- 
gar which sublimes, and arsenic which remains behind. 

It is soluble in yellow ammonium sulphide, forming 
a solution from which arsenic trisulphide is precipi- 
tated on acidification. 

Arsenic disulphide or realgar: Formula — As 2 S 2 . 

This compound occurs native as realgar, crystalliz- 
ing in oblique prisms. These possess an orange-yellow 
color and resinous lustre, and are more or less translu- 
cent, the streak varies from orange-yellow to red. It 
occurs together with silver and lead ores at Andreas- 
berg in the Harz and other localities, and embedded 
in dolomite on the St. Gothard, and has been found 
in minute crystals in Vesuvian lavas. 

When it is heated to 150° with a solution of sodium 
bicarbonate it dissolves in the liquid and afterwards 
separates out in the crystalline form. 

The red arsenic glass or ruby sulphur which occurs 
in commerce is an artificial disulphide of arsenic pre- 
pared in various arsenic works. 

Ruby sulphur is a red glassy mass, translucent at 
the edges. It does not possess a constant composition. 
The material manufactured at Freiberg contains gen- 
erally 75 per cent of arsenic and 25 per cent of sul- 
phur, and that made at Reichenstein in Silesia is a 
mixture of 95 parts of disulphide with 5 parts of sul- 
phur. This body was formerly much used as a pig- 
ment, and is still employed in the manufacture of the 
so-called Indian or white-fire, which is a mixture of 



64 ELEMENTARY CHEMISTRY 

two parts of the disulphide with twenty-four parts of 
nitre, and burns with a splendid white light when 
ignited. The disulphide is also employed in tanning, 
being mixed with lime and employed for removing the 
hair from the skins. 

Arsenic Trisulphide: Formula — As 2 S 3 . 

This substance occurs in nature and is known under 
the name of orpiment or the yellow sulphide of ar- 
senic. It crystallizes in translucent lemon-colored 
prisms. 

When sulphuretted hydrogen is passed through an 
aqueous solution of arsenious oxide the liquid becomes 
of a yellow color, but no precipitate is formed, the 
liquid containing arsenic trisulphide in the colloidal 
form. If a small quantity of hydrochloric acid be 
present, a beautiful yellow precipitate of arsenic tri- 
sulphide is at once thrown down. On heating this sub- 
stance it melts to a yellowish-red liquid which vola- 
tilizes without decomposition at a temperature of about 
700°. Heated in the air it takes fire and burns with 
a pale blue colored flame to arsenious oxide and sul- 
phur dioxide. 

It dissolves in solutions of the alkali hydroxides, a 
brown precipitate of arsenic, containing a little sul- 
phur, separating out slowly, whilst the solution con- 
tains thio-arsenate and thio-oxyarsenates. 

The sulphide of arsenic occurring in commerce is 
prepared by subliming a mixture of seven parts of 
pulverized arsenious oxide with one part of sulphur, 
and it is really a mixture of arsenious oxide with more 
or less sulphide of arsenic. The material thus pre- 
pared, which is very poisonous, from the excess of 



BARIUM 65 

arsenious oxide which it contains, was formerly much 
used as a pigment under the name of King's yellow, 
but it is now almost entirely superseded by the com- 
paratively innocuous chrome yellow. The yellow sul- 
phide of arsenic is also used in the arts and manufac- 
tures, for instance in the printing of indigo colors, and 
a mixture of orpiment, water, and slaked lime is used 
in the East under the name of Rusma as a depilatory, 
its action depending upon the formation of a hydro- 
sulphide of calcium. 

BARIUM. 

Symbol— Ba. Atomic Weight— 136.41. 

Barium.compounds occur somewhat more widely dis- 
persed than those of strontium, the two most common 
barium minerals being the sulphate, or heavy spar, 
and the carbonate. The metal barium has not yet been 
obtained in the coherent state. 

Barium and Chlorine Compounds. 
Chloride of Barium: Formula — BaCl 2 . 

Barium Chloride is one of the most important com- 
pounds of barium, it crystallizes in flat scales contain- 
ing two atoms of water. It may be prepared by dis- 
solving the native carbonate in hydrochloric acid, and 
it is largely used as a precipitant for sulphuric acid. 

Sulphate of Barium: Formula — BaS0 4 . 

Barium Sulphate occurs native and crystalline as 
heavy spar. It is one of the most insoluble salts known, 
and falls as a white crystalline precipitate when any 
soluble barium salt is brought into a solution of a 
sulphate. It is used as a paint, and the precipitated 



66 ELEMENTARY CHEMISTRY 

salt is termed blane fixe, whilst the native heavy spar, 
when ground, is largely used to adulterate white lead. 

Barium and Oxygen Compounds. 
Barium Monoxide or Baryta, BaO and Barium Diox- 
ide, Ba0 2 . 

Barium Monoxide: Formula — BaO. 

The best way of forming this oxide is to decompose 
the nitrate by heat. It is a greyish porous mass, which 
fuses at a high temperature, and takes up water with 
evolution of much heat, forming a crystalline hydrate. 
This hydrate is soluble in twenty parts of cold water, 
and the solution on exposure to the air rapidly absorbs 
carbonic acid, and becomes milky. 

Barium Dioxide: Formula — Ba0 2 . 

When baryta is gently heated in a current of oxy- 
gen gas, the two substances combine together to form 
a dioxide containing twice as much oxygen as baryta, 
this additional atom of oxygen is, however, evolved at 
a higher temperature, and it has been proposed to use 
this decomposition for the manufacture of oxygen from 
the air. For this purpose, as soon as the dioxide Ba0 2 
has been reduced to BaO, the temperature is lowered, 
and air passed over the baryta, this again takes up 
oxygen, passing into Ba0 2 , which again is decomposed 
by a higher temperature. This interesting process has, 
however, been found not to work in practice. There 
are no salts known corresponding to this oxide. 

BISMUTH. 

Symbol— Bi. Atomic Weight— 207.0. 

This metal is found in small quantities in the native 
state, but occurs more often as a sulphide, it is easily 



BISMUTH 67 

reduced to the metallic state, and then exhibits a pink- 
ish-white color. Bismuth does not oxidize in dry air 
at the ordinary temperature, but if heated strongly 
it burns with a blue flame, forming an oxide, it also 
takes fire when thrown into chlorine gas. Bismuth 
dissolves easily in nitric acid. The metal is chiefly used 
as an ingredient of fusible metal, its compounds are 
also used in medicine and as pigments. 

Bismuth and Oxygen Compounds. 

Two oxides of bismuth are known, Bismuth trioxide, 
Bi 2 3 , and Bismuth pentoxide, Bi 2 5 . The first of 
these is a pale yellow powder, formed when the metal 
is roasted in the air, the second oxide is obtained by 
dissolving the first in potash, and precipitating the 
pentoxide by nitric acid and heating, it is a reddish- 
brown powder. Like the corresponding antimony com- 
pound, bismuth pentoxide forms with the alkalies solu- 
ble salts. 

Bismuth and Nitrogen Compounds. 

Bismuth nitrate, Bi(N0 3 ) 3 +5H 2 0, is the most im- 
portant soluble salt of bismuth, the sulphide, Bi 2 S 3 , 
is a black insoluble compound, the trichloride, BiCl 3 , 
is obtained by heating the metal in chlorine. One of 
the most striking peculiarities of the bismuth com- 
pounds is, that solutions of the salts become milky on 
the addition of water, owing to the formation of in- 
soluble basic compounds. Thus Bi(OH) 2 N0 3 is formed 
as a white powder, used in medicine, by adding water 
to a solution of the normal nitrate, and an oxichloride 
BiOCl is precipitated by adding water to the tri- 
chloride. Metallic bismuth is easily reduced from its 
compounds, before the blowpipe, as a brittle bead. 



68 ELEMENTARY CHEMISTRY 

BORON. 

Symbol— B, Atomic Weight— 10.91* 

Boron combined with oxygen and sodium is found 
as borax in nature, it is also found combined with oxy- 
gen alone as boron trioxide. It exists in two forms, 
crystalline and amorphous. Boron is easily obtained 
as a grey amorphous powder, by heating fused boron 
trioxide, B 2 a , with sodium. Crystallized boron is pre- 
pared by heating the amorphous form strongly with 
aluminium, this metal in the fused state having the 
property of dissolving boron, which separates out in 
nearly colorless crystals when the metal cools, just as 
the graphitoidal form of carbon does from its solution 
in iron on cooling. In one specimen of these colorless 
crystals which was analyzed some quantity of carbon 
was found to be present, hence carbon may be said to 
have been prepared artificially in the diamond modifi- 
cation. Boron burns when strongly heated in oxygen 
01 in chlorine, forming the oxide or chloride. It is re- 
markable as being one of the few elements which unite 
directly with nitrogen. 

Amorphous Boron. This modification is obtained by 
heating boron trioxide, obtained by the ignition of 
boric acid, with potassium in an iron tube. It may 
also be obtained by mixing ten parts of coarsely pow- 
dered boron trioxide with six parts of sodium, bringing 
the mixture into a crucible already heated to redness, 
and covering it with a layer of powdered sodium chlo- 
ride previously well dried. As soon as the reaction, 
which is very violent, has subsided, the mass is stirred 
with an iron rod until all the sodium has been oxidized, 
and then carefully poured into water acidified with hy- 



BORON 69 

drochloric acid. The soluble salts dissolve in the water, 
whilst the boron remains behind as an insoluble brown 
powder. This is then collected on a filter, and it must 
be very carefully dried, as it is easily oxidized and 
may take fire. 

The boron obtained in this way is impure, and a 
similar product is got when boron trioxide or borax is 
heated with magnesium powder. In this case a boride 
of magnesium is probably formed which is decom- 
posed by acids. The product left after this treatment, 
however, still contains about 5 per cent of magnesium 
and about 1 per cent of hydrogen, which is also found 
in boron prepared by means of sodium, and is prob- 
ably present in the form of a solid hydride. 

In order to prepare pure boron, 70 grams of mag- 
nesium powder are heated with 210 grams of boron 
trioxide. The product, which consists of boron, accom- 
panied by magnesium boride and borate, is treated 
with dilute acid which dissolves the borate and the 
greater part of the boride. In order to remove the 
last portions of the boride the residue is then fused 
with borax and again treated with hydrochloric acid, 
which leaves the boron containing only traces of sili- 
con, iron, and magnesium. The presence of nitride of 
boron in the final product can only be avoided by car- 
rying out the heating in an atmosphere of hydrogen or 
by placing the mixture in a crucible lined with titanic 
acid. 

Amorphous boron is a chestnut-brown powder. It 
is infusible even at the temperature of the electric arc, 
but volatilizes, the extremities of the electrodes being 
converted into boron carbide, while if it be heated to 
700° in air it burns. It is a more powerful reducing 



70 ELEMENTARY CHEMISTRY 

agent than carbon or silicon, inasmuch as it is oxidized 
by carbonic oxide and by silica. It combines with 
bromine at 700°, but not with iodine, and forms com- 
pounds with many metals on heating, among which are 
silver and platinum, while on fusing with cast-iron in 
a current of hydrogen it replaces the carbon giving a 
white, hard cast-iron, and if fused with reduced iron 
produces brown steel. It is acted on by the oxy-acids, 
is oxidized by water vapor, and combines with nitro- 
gen at a high temperature. It is a non-conductor of 
electricity, and when freshly prepared and not strong- 
ly ignited is slightly soluble in water, imparting to it 
a yellow color and being precipitated unchanged from 
its aqueous solution on the addition of acids or salts. 

Crystalline or Adamantine Boron. This substance 
was first obtained in the year 1856. It can be prepared 
by several processes. Thus, if amorphous boron be 
pressed down tightly in a crucible, a hole bored in the 
center of the pressed mass, a rod of aluminium dropped 
into the hole, and the crucible then heated to white- 
ness, the boron dissolves in the molten aluminium and 
separates out in the crystalline form when the metal 
cools. The aluminium is then dissolved in caustic 
soda, and thus the insoluble boron is left in large trans- 
parent yellow or brownish-yellow crystals. The same 
modification may be obtained in smaller crystals, which 
are often joined together in the form of long prismatic 
needles, by melting together boron trioxide and alumin- 
ium. In order to prevent the action of the oxygen of 
the air upon the fused mass, the crucible in which the 
operation is conducted must be placed inside a larger 
one and the space between them filled up with pow- 
dered charcoal. In this process, however, the boron 



BORON 71 

takes up carbon to the amount of from 2 to 4 per cent. 
This carbon must be in the form of diamond carbon, 
inasmuch as the boron crystals containing this im- 
purity are transparent, and more transparent the 
larger the percentage of carbon. In addition to car- 
bon the boron thus prepared is found to contain a cer- 
tain quantity of iron and silicon from the crucible used. 
These impurities can be removed by treatment with 
hydrochloric acid, and afterwards with a mixture of 
nitric and hydrofluoric acids. The crystals of adaman- 
tine boron contain aluminium as well as carbon, and 
possess a constant composition. 

It may also be prepared by heating boric acid with 
aluminium turnings and sulphur, and treating the 
cooled mass with water, when aluminium sulphide is 
decomposed and crystalline boron left. 

Boron crystallizes in monoclinic pyramids or prisms, 
which have a lustre and a hardness exceeded only by 
that of the diamond, as they scratch both ruby and 
corundum. When heated in the air or in oxygen it 
ignites at the same temperature as does the diamond, 
and then does not oxidize throughout the mass, but be- 
comes covered with a coating of the melted trioxide. 
Concentrated nitric acid exerts no action upon it, and 
even aqua regia attacks it but slowly. Boiling caustic 
soda solution likewise does not act upon it, but if it is 
fused with the solid alkali it dissolves slowly with for- 
mation of sodium borate and with evolution of hydro- 
gen. 

Boron and Carbon Compounds. 
Boron Carbide: Formula — B 2 C 2 . 

This is prepared by heating a mixture of boric an- 
hydride and carbon in an electric furnace, and forms 



72 ELEMENTARY CHEMISTRY 

a graphite-like powder which melts at a very high 
temperature, burns with difficulty in oxygen, is in- 
soluble in the usual solvents, but is decomposed by 
fusion with alkalies. 

Another boride of carbon, having the composition 
CB 6 , is prepared by heating sixty-six parts of amor- 
phous boron and twelve parts of carbon, obtained from 
sugar, in an electric furnace, or by dissolving the two 
elements in iron, copper, or silver at a very high tem- 
perature, and removing the metal by treatment with 
aqua regia. It forms very hard black lustrous crystals, 
having a density of 2.41, which are attacked by 
chlorine below 1,000°, and by oxygen very slowly at 
that temperature, like the previous compound, it is 
decomposed by fusion with alkalies. The powder is 
sufficiently hard to cut the diamond but more slowly 
than diamond dust. 

BROMINE. 
Symbol— Br. Atomic Weight— 79.36. 

This element, which closely resembles chlorine in 
its properties and compounds, was discovered in the 
salts obtained by the evaporation of sea-water. It does 
not occur free in nature, and is, like chlorine, found 
combined with sodium and magnesium in the waters 
of certain mineral springs. In order to obtain pure 
bromine, use is made of the fact that free chlorine 
liberates bromine from its combinations with metals, 
forming a metallic chloride. The bromine thus set 
free may be separated by shaking the liquid up with 
ether, which dissolves the bromine, forming a bright 
red solution. On adding caustic potash to this ethereal 
solution, the color at once disappears, the bromine be- 






BROMINE 73 

comes combined, forming the bromide and bromate of 
potassium, on evaporation of the ether these salts re- 
main, and after ignition to decompose the bromate, the 
bromide can again be liberated by the action of sul- 
phuric acid and manganese dioxide, exactly as in the 
case of chlorine. 

One part of bromine dissolves in about 30 parts of 
water at 15°, and this solution possesses bleaching 
powers, feebler, however, in action than those of chlo- 
rine. This bleaching action is caused by the oxidation 
of the coloring matter, the bromine combining with 
the hydrogen of the water to form an acid called hy- 
drobromic acid, corresponding in mode of formation 
and properties to hydrochloric acid. 

In order to detect bromine in a mineral water, or to 
prepare it in small quantities, the following method 
is employed. The mother-liquor remaining after the 
brine has been well crystallized is treated with a 
stream of chlorine gas, so long as the yellow color of 
the liquid continues to increase in depth. Chlorine has 
the power of liberating bromine from bromides, itself 
uniting with the metal, and the bromine being set free, 
thus : 

MgBr 2 +Cl 2 =MgCl 2 +Br 2 . 
The addition of excess of chlorine is to be avoided, as 
a compound of chlorine and bromine is then formed. 
The yellow liquid is then well shaken with chloroform, 
which dissolves the bromine, forming, on standing, a 
brown solution below the aqueous liquid. On adding 
caustic potash to this solution the color at once disap- 
pears, the bromine combining to form the bromide, 
KBr, and bromate of potassium, KBr0 3 , thus : 
3Br 2 +6KHO=KBr0 3 +5KBr+3H 2 0. 



74 ELEMENTARY CHEMISTRY 

On further concentrating the solution a mixture of 
these salts remains, and from these the bromine is 
again liberated by distilling the liquid with black 
oxide of manganese and sulphuric acid in a tubulated 
retort. The decomposition which here occurs is simi- 
lar to that which takes place in the preparation of 
chlorine. Dark red fumes of bromine are liberated, 
and a black liquid condenses in the well-cooled re- 
ceiver. 

If the bromine is required to be anhydrous it must 
be re-distilled over concentrated sulphuric acid, and 
if iodine is present this must be got rid of previously 
by precipitation as subiodide of copper. 

By far the greater quantity of the bromine brought 
into commerce is now manufactured at Stassfurt from 
the mother-liquor remaining after the separation of 
the potassium salts contained in the salt deposits, the 
process adopted consisting in the treatment of the 
liquors with chlorine under suitable conditions. 

Bromine is a heavy mobile liquid, so dark as to be 
opaque except in thin layers. It is the only liquid ele- 
ment at the ordinary temperature except mercury. It 
freezes at — 7° to a dark brown solid, evaporates quick- 
ly in the air, boils at 59°, and crystallizes from carbon 
bisulphide at — 90° in slender dark carmine-red 
prisms. Bromine possesses a very strong, unpleasant 
smell, the vapors when inhaled produce great irrita- 
tion, and affect the eyes very painfully. When swal- 
lowed it acts as an irritant poison, and when dropped 
on the skin it produces a corrosive sore which is very 
difficult to heal. 

In its general properties, as well as in those of its 
compounds, bromine closely resembles chlorine, al- 



BROMINE 75 

though they are not so strongly marked. Thus it 
bleaches organic coloring matters, but much less quick- 
ly than chlorine, and it combines directly with metals 
to form bromides, though its action is less energetic 
than that of chlorine. It does not combine at all at 
ordinary temperatures with metallic sodium, indeed 
these two substances may be heated together to 200° 
before any perceptible action commences, whereas bro- 
mine and potassium cannot be brought together with- 
out combination occurring, sometimes with almost ex- 
plosive violence. The addition, however, of a drop of 
water to bromine and clear sodium sets up a lively 
reaction. 

The solution of bromine in water has an orange-red 
color, it soon loses bromine in contact with the air, and 
bleaches organic coloring matter. Bromine water is 
permanent in the dark, but on exposure to sunlight it 
becomes acid from the formation of hydrobromic acid 
and evolution of oxygen. Bromine also dissolves read- 
ily in chloroform, carbon bisulphide, alcohol, ether, and 
acetic acid. 

Bromine and Hydrogen Compounds. 
Hydrobromic Acid: Formula — HBr. 

Hydrobromic acid, or hydrogen bromide, is prepared 
by the action of acids (phosphoric acid) on the bro- 
mides, or better by bringing bromine and phosphorus 
in contact with water, when a violent action occurs, 
hydrobromic acid and phosphoric acid being formed 
thus : 

P + 5Br+4H 2 0=5HBr+H 3 P0 4 . 

It is colorless gas, having a strong acid reaction, and 
fumes strongly in moist air, it is very soluble in water. 
When concentrated, the aqueous acid boils (under 760 



76 ELEMENTARY CHEMISTRY 

mm. pressure) at 126°, and contains 47.8 per cent of 
HBr. Two volumes of this gas contain one of bromine 
united with one of hydrogen. The aqueous acid neu- 
tralizes bases, forming the bromides and water. The 
gas liquefies at — 73°. 

CADMIUM. 
Symbol— Od. Atomic Weight— 111.3. 

This is a comparatively rare metal, occurring in 
small quantities in most zinc ores. In its chemical re- 
lations it closely resembles zinc. It is, however, more 
volatile than the latter metal, and therefore distills over 
first in the preparation of zinc. Cadmium is a white 
ductile metal, melting at 315°, it may be easily distin- 
guished and separated from zinc by yielding a bright 
yellow sulphide which is insoluble in hydrochloric acid. 

CALCIUM. 
Symbol— Ca. Atomic Weight— 39.7. 

Calcium forms a considerable portion of the plutonic 
rocks of which the earth is composed, and occurs in 
very large quantities, forming whole mountain-chains 
of limestone, chalk, gypsum, and mountain limestone. 
The metal calcium is obtained by the decomposition of 
the chloride by the electric current, or by heating the 
iodide with sodium, it is a light yellow metal which 
easily oxidizes in the air, and when heated in air it 
burns with a bright light, lime, the only oxide of cal- 
cium, being formed. 

Calcium and Chlorine Compounds. 
Calcium Chloride: Formula — CaCl 2 . 

This soluble salt is formed when limestone or mar- 
ble is dissolved in hydrochloric acid. If the solution 



CALCIUM 77 

be then evaporated, colorless needle-shaped crystals of 
the hydrated chloride are formed. When these are 
dried, the substance still retains 2H 2 0, and forms a 
porous mass which takes up moisture with great avid- 
ity, and is much used for drying gases. When this 
mass is more strongly heated, it fuses and parts with 
all its water. 

Calcium and Oxygen Compounds. 
Calcium Oxide or Lime: Formula — CaO. 

Pure lime is obtained by heating white or black mar- 
ble to redness in a vessel exposed to the air. Lime is 
prepared on a large scale for building and other pur- 
poses, by heating limestone in kilns by means of coal 
mixed with the stone, the carbonic acid escapes, and 
quick or caustic-lime remains. Pure lime is a white 
infusible substance, which combines with water very 
readily, giving off great heat, and falling to a white 
powder called calcium hydroxide, or slaked lime. The 
hydrate is slightly soluble in water, 1 part of it dis- 
solving in 730 parts of cold, but only in 1300 parts of 
boiling water, and forming lime-water, which, like the 
hydrate, has a great power of absorbing carbonic acid 
from the air. It is indeed owing to this property that 
the hardening or setting of mortars and cements made 
from lime is due. Mortar consists of a mixture of 
slaked lime and sand. A gradual combination of the 
lime with the silica occurs, and this helps to harden 
the mixture. Hydraulic cements, which harden under 
water, are prepared by carefully heating an impure 
lime containing clay and silica. A compound silicate 
of lime and alumina appears to be formed on moisten- 
ing the powder, which then solidifies, and is unacted 



78 ELEMENTARY CHEMISTRY 

upon by water. Lime is largely used in agriculture, 
its action being to destroy the excess of vegetable mat- 
ter contained in the soil, and to liberate the potash for 
the use of the plants from heavy clay soils by decom- 
posing the silicate. 

Calcium, Carbon and Oxygen Compounds. 
Calcium Carbonate: Formula — CaC0 3 . 

This salt occurs most widely diffused, as chalk, lime- 
stone, coral, and marble, many of those enormous de- 
posits being made up of the microscopic remains of 
minute sea-animals. Calcium carbonate exists crys- 
talline as calc-spar, or Iceland spar, and also in a dif- 
ferent form, so that this substance is dimorphous. The 
carbonate is almost insoluble in pure water, but read- 
ily dissolves when the water contains carbonic acid, 
giving rise to what is termed temporarily hard water. 
Such a water deposits a crust of calcium carbonate on 
boiling, owing to the escape of the carbonic acid. The 
well-known evil of boiler incrustation is caused by 
these deposits. The formation of such a crust may be 
checked, if not avoided, by adding a small quantity 
of sal-ammoniac to the water, soluble calcium chloride 
and volatile ammonium carbonate being formed. 
Water hard with dissolved carbonate may be softened 
by the addition of lime suspended in water in such 
quantity that the excess of carbonic acid is neutralized. 

Calcium, Chlorine and Oxygen Compounds. 
Chloride of Lime: Formula— CaCl 2 . Ca2C10. 

Bleaching Powder, or Chloride of Lime, is a mixture 
of calcium chloride and calcium hypochlorite, and is 



CARBON 79 

obtained by the action of chlorine upon slaked lime. 
If a clear solution of bleaching powder is heated with 
a small quantity of oxide of cobalt or of copper, the 
oxygen of the hypochlorite is gradually evolved, and 
calcium chloride left behind. This decomposition de- 
pends upon the fact that higher oxides of the metal 
are at first formed, but these decompose under the in- 
fluence of heat, and give off oxygen, regenerating the 
lower oxide, which again attacks another portion of 
hypochlorite, and thus the process becomes continu- 
ous. It is not improbable that the action of manga- 
nese dioxide in facilitating the evolution of oxygen 
from potassium chlorate may depend upon a similar 
action. 

Calcium Sulphate : Formula — CaS0 4 . 

This occurs in nature as a mineral termed Anhy- 
drite, and combined with 2H 2 as selenite, gypsum, or 
alabaster. It is soluble in 400 parts of water, and is 
a very common impurity in spring water, giving rise 
to what is termed permanent hardness, as it cannot be 
removed by boiling. Gypsum when moderately heated 
loses its water, and is then called plaster of Paris, this 
when moistened takes up two atoms of water again 
and sets to a solid mass, and is therefore much used 
for making casts and moulds. 

CARBON. 

Carbon: Symbol — C. Atomic Weight — 16. 

Carbon is remarkable as existing in three distinct 
forms, which, in outward appearance or physical prop- 
erties, have nothing in common, whilst their chemical 
relations are identical. These three allotropic forms 



80 ELEMENTARY CHEMISTRY 

of carbon are Diamond, Graphite or Plumbago and 
Charcoal. These substances differ in hardness, color, 
specific gravity, etc., but they each yield on combus- 
tion in the air or oxygen the same weight of the same 
substance, carbonic acid, or carbon dioxide, 12 parts 
by weight of each of tliese forms of carbon yielding 
44 parts by weight of carbon dioxide. Carbon is the 
element which is specially characteristic of animal and 
vegetable life, as every organized structure, from the 
simplest to the most complicated, contains carbon. If 
carbon were not present on the earth, no single veg- 
etable or animal body such as we know could exist. 
In addition to the carbon which is found free in these 
three forms, and that contained combined with hydro- 
gen and oxygen in the bodies of plants and animals, 
it exists combined with oxygen as free carbon dioxide 
in the air, and with calcium and oxygen as calcium 
carbonate in limestone, chalk, marble, corals and shells. 
The fact has already been noticed that plants are able, 
when exposed to sunlight, to decompose the carbon 
dioxide in the air, liberating the oxygen, and taking 
the carbon for the formation of their vegetable struc- 
ture, whilst all animals, living directly or indirectly 
upon vegetables, absorb oxygen, and evolve carbon di- 
oxide. Thus the sun's rays, through the medium of 
plants, effect deoxidation or reduction, while animals 
act as oxidizing agents with respect to carbon. 

The element carbon not only combines directly with 
oxygen, but also with hydrogen, forming a compound 
called acetylene, C 2 H 2 . Carbon forms with oxygen, 
hydrogen, and nitrogen a series of more or less compli- 
cated compounds very much more extended than the 
series formed with these bodies by any other element, 



CARBON 81 

so that these compounds are considered as a separate 
branch of the science under the name of Organic Chem- 
istry. 

The Diamond was first found to consist of pure car- 
bon by Lavoisier, in 1775-6, by burning it in oxygen, 
and collecting the carbon dioxide formed, it occurs 
crystallized in certain sedimentary rocks and gravel in 
India, Borneo, and Brazil. Diamond occurs crystal- 
lized in forms derived by a symmetrical geometric op- 
eration from a regular octahedron. The specific grav- 
ity of diamond varies from 3.3 to 3.5, it is the hardest 
of all known bodies, and when cut possesses a brilliant 
lustre, and a high refractive power. In addition to its 
employment as a gem, the diamond is used for cutting 
and writing upon glass. We are altogether unacquaint- 
ed with the mode in which the diamond has been 
formed, it cannot, however, have been produced at a 
high temperature, because, when heated strongly in 
a medium incapable of acting chemically upon it, the 
diamond swells up, and is converted into a black mass 
resembling coke. 

Graphite, or Plumbago, crystallizes in six-sided 
plates which have no relation to the form in which the 
diamond crystallizes. Graphite occurs in the oldest 
sedimentary formations, and in granitic or primitive 
rocks, it is found in large quantities in Siberia and 
Ceylon. It has a black metallic appearance, whence 
the familiar name black lead, and leaves a mark when 
drawn upon paper. The specific gravity of graphite 
is 2.15 to 2.35. Coarse impure graphite maybe puri- 
fied by heating the powder with sulphuric acid and 
potassium chlorate, a compound is thus obtained which, 
on being heated strongly, decomposes, leaving pure 



82 ELEMENTARY CHEMISTRY 

graphite in a bulky and finely-divided powder. This 
powder when strongly compressed forms a coherent 
mass, from which pencils and other articles can be 
made. Graphite is used for polishing surfaces of iron- 
work, and also for giving a protecting varnish to 
grains of gunpowder. Graphite is produced in the 
manufacture of iron, it occasionally separates from 
the molten pig-iron in the form of scales. 

Charcoal is the third allotropic modification of car- 
bon. It is obtained in a more or less pure state when- 
ever animal or vegetable matter is heated to redness 
in a vessel nearly closed, the volatile matters, com- 
pounds of carbon, hydrogen, and oxygen, are thus 
driven off, and the residue of the carbon, together with 
the ash or mineral portion of the organism, remains 
behind. 

The purest form of charcoal-carbon is found in lamp- 
black, it also occurs as wood charcoal, coal, coke, and 
animal charcoal. This form of carbon does not crys- 
tallize, and is hence termed amorphous carbon. It is 
much lighter than either of the other two forms, the 
specific gravity of powdered coke varying from 1.6 to 
2.0. Charcoal appears at first sight to be lighter than 
water, as a piece of it floats on the surface of this 
liquid, this is, however, due to the porous nature of 
the charcoal, for if it be finely powdered it sinks to 
the bottom of the water. This porous nature of char- 
coal enables it to exert a remarkable absorptive pow- 
er, of which much use is made in the arts. Charcoal 
is thus able to absorb about ninety times its own vol- 
ume of ammonia gas, and about nine volumes of oxy- 
gen. In the process of sugar-refining, use is made of 
the property of charcoal to absorb the coloring mat- 



CARBON 83 

ters present in the raw sugar. The kind of charcoal 
best suited to this purpose is that obtained by heating 
bones in a closed vessel. Charcoal is also used as a 
disinfectant in hospitals and dissecting rooms. It ap- 
pears that the putrefactive gases when absorbed by 
the charcoal undergo a gradual oxidation from contact 
with the oxygen of the air taken up by the charcoal, 
and are thus rendered harmless. 

Coal is a form of carbon less pure than wood char- 
coal. It consists of the remains of a vegetable world 
which once flourished on the earth's surface. The 
original woody fibre has undergone a remarkable 
transformation in passing into coal, having been sub- 
jected to a process similar, in a chemical point of view, 
to that by which wood is transformed into charcoal. 
It has not, however, lost the whole of its hydrogen and 
oxygen, and it has at the same time become bitumen- 
ized, so that for the most part all the vegetable struc- 
ture has disappeared. There are many different kinds 
of coal, containing more or less of the oxygen and hy- 
drogen of the original wood. Cannel coal and bog coal 
contain the most hydrogen, and anthracite coal the 
least. 

Carbon and Hydrogen Compounds. 

Methane or Carburetted Hydrogen, CH 4 , Acetylene, 
C 2 H 2 and Ethylene or defiant Gas, C 2 H 4 . 
These compounds are very numerous, they are 
known in the gaseous, liquid, and solid forms. A still 
larger number of substances exist containing carbon, 
hydrogen, and oxygen, with sometimes nitrogen, these 
are termed organic compounds, and they are more 
numerous than all the compounds of the other elements 



84 ELEMENTARY CHEMISTRY 

put together. Many of these are found to be formed 
from the bodies of plants and animals. 

Methane or Marsh Gas: Formula— CH 4 . Molecular 
Weight— 15.9. 

This is a colorless, tasteless, inodorous gas, which 
has not been condensed to a liquid. It is found in coal 
mines, and known under the name of firedamp, it also 
occurs in stagnant pools, being produced by the decom- 
position of dead leaves whence the name marsh gas. 
It is one of the constituents of coal gas, etc., and is 
evolved in many volcanic districts. Marsh gas may 
also be artificially prepared by heating sodium acetate 
with caustic soda. 

Sodium acetate and caustic soda give sodium car- 
bonate and marsh gas. 

Marsh gas burns with a bluish-yellow, non-luminous 
flame, forming carbon dioxide and water; with a lim- 
ited supply of air it yields several products, amongst 
which is acetylene. If mixed with ten times its volume 
of air, or twice its volume of oxygen, it ignites with a 
sudden and violent explosion on the application of a 
light, and hence the great damage produced by the 
escape of this gas in coal mines. The composition of 
marsh gas is ascertained by exploding it with oxygen 
in the eudiometer. 1 volume of this gas and 3 volumes 
of oxygen yield 2 volumes after passage of the spark. 
On absorbing by potash the carbon dioxide produced, 
1 volume of oxygen is found to remain. Hence of 
the 2 volumes of oxygen needed to burn the 1 volume 
of marsh gas, 1 has gone to unite with the carbon, and 

1 to form water with the hydrogen. It is thus seen that 

2 volumes of marsh gas contain 4 volumes of hydrogen 



CARBON 85 

weighing 4, as water contains 2 volumes of hydrogen 
and 1 of oxygen, and as much carbon as is contained 
in 2 volumes of carbon dioxide, viz., 12 parts by 
weight, and hence the formula CH 4 is given to this 
gas. 

Acetylene : Formula — 2 H 2 . 

This gas is formed by the direct union of carbon and 
hydrogen at a very high temperature. For this pur- 
pose the carbon terminals of a powerful galvanic bat- 
tery are brought together in an atmosphere of hydro- 
gen. At the high temperature thus evolved, a direct 
union of carbon and hydrogen takes place, and acety- 
lene is formed. Acetylene is a colorless gas, which 
burns with a bright luminous flame, and possesses a 
disagreeable and very peculiar odor, it is produced in 
all cases of incomplete combustion, and its smell may 
be noticed when a candle burns with a smoky flame. 
Acetylene combines with certain metals, such as cop- 
per and silver, and the compounds thus formed are 
distinguished by the ease with which they undergo 
explosive decomposition. This gas likewise unites di- 
rectly with hydrogen, forming the next substance, 
ethylene, C 2 H 2 +H 2 =C 2 H 4 . 

Acetylene is now manufactured on a large scale for 
use as an illuminant in places where coal-gas is not 
available, such as small villages and country houses 
or railway stations. The method invariably adopted 
for its production is the action of water on calcium 
carbide. The carbide was, however, first obtained in 
quantity in 1892 by a Canadian engineer named Will- 
son, in an attempt to manufacture metallic calcium 
by heating lime with carbonaceous matter in the elec- 



86 ELEMENTARY CHEMISTRY 

trie furnace, the resulting product being, however, cal- 
cium carbide. 

The manufacture of acetylene from carbide was 
commenced about 1895, and since that date some hun- 
dreds of different forms of apparatus have been pat- 
ented for carrying out the decomposition of the car- 
bide with water. These may all be divided into two 
general classes : Those in which the water is added to 
the carbide, and those in which the carbide is added 
to the water so that the latter is always in excess. The 
objection to generators of the first class is the higher 
temperature attained in the reaction, which results in 
the partial conversion of acetylene into tarry matters. 
Some generators of each class are constructed so as 
to render them "automatic," the generation of acety- 
lene only takes place as fast as it is consumed, whilst 
in others the whole charge is decomposed at once, and 
the resulting gas stored in a holder for use as required. 

Ethylene or Olefiant Gas: Formula— C 2 H 4 . 

This gas is obtained on the destructive distillation of 
coal, and is an important constituent of coal gas. It 
is obtained in the pure state by heating 1 part of 
alcohol (spirits of wine), C 2 H 6 0, with 5 or 6 parts by 
weight of strong sulphuric acid, as in formation of 
carbon monoxide from formic acid the elements of 
water are separated by the sulphuric acid, and C 2 H 4 
is evolved as a gas. This gas is colorless, but possesses 
a sweetish taste, by exposure to a high pressure at 
a temperature of — 110° it has been condensed to a 
colorless liquid. On bringing it in contact with a 
light in the air, it burns with a luminous smoky flame, 
forming carbon dioxide and water. When mixed 



CARBON 87 

with three times its bulk of oxygen and fired, it de- 
tonates very powerfully. 1 volume of olefiant gas re- 
quires 3 volumes of oxygen to burn it completely, and 
yields 2 volumes of carbon dioxide, so that 1 volume of 
oxygen is needed to combine with the hydrogen. Hence 
this gas contains twice as much carbon as marsh gas, 
with the same quantity of hydrogen, we must there- 
fore write its formula C 2 H 4 . 

Olefiant gas combines directly with its own volume 
of chlorine gas, forming an oily liquid, C 2 H 4 C1 2 ; and 
owing to this property it has received the above name. 

Carbon, Hydrogen and Oxygen Compounds. 
Ether or Diethyl Ether: Formula— C 4 H 10 O. 

Ether is prepared on a large scale by heating a 
mixture of alcohol and sulphuric acid to 140°, when 
ether and water are given off. The decompositions 
which take place are as follows: In the first place, 
alcohol and sulphuric acid form hydrogen-ethyl-sul- 
phate (sulphovinic acid) and water, by an exchange of 
hydrogen and ethyl. 

This hydrogen-ethyl-sulphate next comes in contact 
with a second molecule of alcohol, another exchange 
of hydrogen for ethyl occurs, and ether and sulphuric 
acid are formed. 

The water formed by the first decomposition, and 
the ether produced by the second, are given off as 
vapor, whilst the sulphuric acid remains behind, ready 
again to go through the same series of changes on 
meeting with two other molecules of alcohol. This 
process is called the continuous etherification process, 
as a current of alcohol may be passed continuously 



ELEMENTARY CHEMISTRY 

through the sulphuric acid heated to 140° , and a 
regular supply of ether and water thus obtained. 

Ether is a colorless, very mobile liquid, possessing 
a strong and peculiar smell. It is lighter than water, 
and is not miscible with that liquid. Ether boils at 
34°, and its vapor is 37 times heavier than hydrogen, 
and can be poured from vessel to vessel like carbonic 
acid gas. It burns with a luminous flame, and ex- 
plodes when mixed with air. From its low boiling 
point great care must be taken to avoid explosions 
when working with this substance, owing to the vapor 
becoming mixed with air. Ether is easily attacked 
by oxidizing agents, yielding the same products as 
alcohol, it is also acted upon by chlorine. 

Carbon and Nitrogen Compounds. 
Cyanogen Gas: Formula — C 2 N 2 . 

This gas is formed when the cyanides of mercury, 
silver, or gold are heated. For this purpose it is usual 
to employ mercuric cyanide, which decomposes thus 

Hg(CN) 2 =Hg+C 2 N 2 . 

The mercury salt, placed in a tube of hard glass 
fitted with a cork and gas-delivery tube, or in a small 
hard glass retort, is heated to dull redness, and the 
gas collected over mercury. 

The heat of formation of mercuric cyanide is large, 
and a very high temperature is therefore required to 
bring about this decomposition. If mercuric chloride 
be mixed with the cyanide, the cyanogen comes off 
at a lower temperature, as the mercury then combines 
with the mercuric chloride forming mercurous chlor- 



CAEBON 89 

ide, and the whole reaction, represented by the equa- 
tion 

Hg(CN) 2 +HgCl 2 =(CN) 2 +2HgCL 

It is also formed by gently igniting an intimate 
mixture of two parts of well-dried potassium ferrocy- 
anide, Fe(CN) 6 K 4 , with three parts of mercuric chlor- 
ide. The gas thus prepared contains free nitrogen, 
and the employment of potassium cyanide in place of 
the ferrocyanide is recommended. There is no doubt 
that in this reaction mercuric cyanide is first obtained, 
and this is then decomposed as already described. 

Cyanogen is most readily obtained by gradually 
adding a concentrated solution of potassium cyanide 
to a solution of two parts of crystallized copper sul-< 
phate in four parts of water, and then warming, cy- 
anogen is evolved and cuprous cyanide simultaneously 
separates out from the solution. If the cuprous cyanide 
be washed by decantation, and then treated with 
ferric chloride, the remainder of the cyanogen is 
evolved. 

Cyanogen gas is found in small quantities in the 
gases proceeding from the blast furnaces, and it is 
likewise formed w T hen a mixture of ammonia and coal- 
gas is burnt in a Bunsen burner. The following equa- 
tion probably explains the reaction which takes place 
in the latter case. 

4CO+4NH 3 +0 2 =2C 2 N 2 +6H 2 0. 

Cyanogen is a colorless gas possessing a peculiar 
pungent odor resembling that of peach kernels. It is 
poisonous, and burns, v when ignited in air, with a 
characteristic purple-mantled flame, with formation 
of carbon dioxide and nitrogen. 



90 ELEMENTARY CHEMISTRY 

Carbon, Hydrogen and Chlorine Compounds. 
Chloroform : Formula — CHCL. 

Chloroform is formed when chlorine acts on marsh 
gas, but it is prepared by acting upon methyl or ethyl 
alcohol with bleaching powder. It is a mobile, heavy 
liquid, possessing a powerful and agreeable smell. 
Chloroform is much used in medicine, producing, when 
it is inhaled, a temporary but perfect insensibility to 
pain, and is therefore much valued in surgical opera- 
tions. 

Carbon, Hydrogen and Oxygen Compounds. 
Methyl or Wood Alcohol: Formula— CH 3 .OH. 

Methyl alcohol is produced by the dry distillation 
of wood, forming about one per cent of the aqueous 
distillate. It is likewise met with in the oil of winter- 
green. Methyl alcohol can likewise be synthetically 
built up from its constituent elements by means of 
several complicated reactions. 

Pure methyl alcohol is obtained from crude wood 
alcohol, in which it is contained mixed with a variety 
of other organic compounds, by forming a crystalline 
methyl oxalate. This, on treatment with water, is 
decomposed, and yields the alcohol in the pure state. 
Methyl alcohol is a colorless, mobile liquid, posses- 
sing a pure spirituous smell. It burns with a non- 
luminous flame, and is soluble in and miscible with 
water. Potassium dissolves in methyl alcohol with 
evolution of hydrogen and formation of potassium 
methylate. Methyl alcohol when acted on by oxidizing 
agents yields methyl aldehyde and formic acid. By 



CAEBON 91 

the action of bleaching-powder on methyl alcohol, 
chloroform is obtained, acted upon by hydrochloric 
acid, the alcohol yields methyl chloride. 

Ethyl or Absolute Alcohol: Formula — C 2 H 5 .OH. 

Ethyl Alcohol is obtained in the vinous fermenta- 
tion of sugar, a decomposition effected in aqueous 
sugar solutions in presence of yeast, in which alcohol 
and carbonic acid are chiefly formed. 

Alcohol and alcoholic liquids are prepared in large 
quantities by the fermentation of sugar derived from 
various sources. The fermented liquid is distilled, 
and the dilute aqueous spirit thus separated from non- 
volatile impurities. It is obtained in a more concen- 
trated form by repeated rectifications, as it boils at 
a lower temperature than water. Alcohol cannot, how- 
ever, be completely separated from water by simple 
distillation, the strongest spirit which can thus be 
prepared containing 10 per cent of water. To with- 
draw all the water, the spirit must be distilled with 
some substance capable of combining with water, such 
as potassium carbonate or quicklime. The pure liquid 
thus obtained is termed absolute alcohol. It is a color- 
less, mobile liquid, possessing a pleasant, spirituous 
smell and burning taste and it boils at 78°. 4 when the 
barometer stands at 760 mms. It has not been solidi- 
fied, becoming only viscid at a temperature of — 100°. 
Alcohol is very inflammable, burning with a slightly 
luminous blue flame. It absorbs moisture with great 
avidity, and mixes with water in all proportions, the 
mixture evolving heat and undergoing contraction in 
volume. 



92 ELEMENTARY CHEMISTRY 



Carbon and Oxygen Compounds. 
Carbon Dioxide, C0 2 and Carbon Monoxide, CO. 

Carbon Dioxide or Carbonic Acid Gas: Formula — 
C0 2 . Molecular Weight— 43.7. 

The carbonic acid contained in the air is derived 
from a variety of sources, it is formed by the respira- 
tion of man and animals, as well as in the act of com- 
bustion of organized material, and in its decay and 
decomposition. The amount of atmospheric carbonic 
acid varies between certain narrow limits, but on an 
average reaches 3 volumes in 10,000 volumes of air. In 
the presence of the sunlight, plants have the power, 
through their leaves, of decomposing this carbonic 
acid, taking up the carbon to form their own tissue, 
and eliminating the oxygen gas, hence the amount of 
carbonic acid in the air does not increase beyond the 
limits named. 

Carbon dioxide is an acid-forming oxide giving rise 
to a series of salts termed the carbonates, many of 
which occur in nature as minerals. Amongst these is 
especially to be mentioned calcium carbonate, CaC0 3 , 
which occurs in two distinct crystalline forms as calc- 
spar and arragonite, while it is found in a massive 
crystalline form in marble and limestone, calcium car- 
bonate also forms the chief constituent of the shells of 
mollusca and foraminifera, the remains of w T hich con- 
stitute the chalk formation as well as the greater part 
of all the limestones. In order to prepare carbon diox- 
ide, a carbonate, such as marble or chalk, is brought 
into a gas-evolution flask, and dilute hydrochloric acid 



CAEBON 93 

poured upon it, when the gas is rapidly evolved with 
effervescence, thus: 

CaC0 3 .+2HCl=C0 2 +CaCl 2 +H 2 0. 

The gas thus obtained invariably carries over small 
quantities of hydrochloric acid vapor with it, from 
which it may be freed by passing through a solution of 
sodium bicarbonate. In order to obtain a constant 
stream of carbon dioxide the same apparatus may be 
employed which is made use of for the preparation 
of sulphuretted hydrogen gas. 

Another method of obtaining a constant current is 
to pour concentrated sulphuric acid over chalk, and 
add a very small quantity of water, thus: 

CaC0 3 +H 2 S0 4 =C0 2 +CaS0 4 +H 2 0. 

The dilute acid cannot be employed for this purpose, 
since the calcium sulphate formed, which is soluble in 
the concentrated acid, does not dissolve in the dilute 
acid, and thus prevents its further action on the chalk. 

The gas obtained from chalk possesses a peculiar 
smell, which is due to the presence of small quantities 
of volatile organic matter always contained in the chalk. 
In order to prepare a very pure gas, sodium carbonate 
may be decomposed with pure dilute sulphuric acid, 
thus : 

Na 2 C0 3 +H 2 S0 4 =C0 2 + Na 2 S0 4 +H 2 0. 

Carbon dioxide may be obtained on a large scale for 
the preparation of sodium bicarbonate, white lead, and 
other commercial products by burning limestone, or 
by the combustion of charcoal or coke, but the gas 
then always contains large volumes of nitrogen de- 
rived from the air used for combustion. In order to 
obtain the dioxide in comparatively pure condition 



94 ELEMENTARY CHEMISTRY 

from such gases on the large scale, the mixture may 
be passed into a concentrated solution of potassium 
carbonate with which carbon dioxide combines forming 
potassium hydrogen carbonate, KHC0 8 . The solution 
of the latter evolves carbon dioxide on warming, leav- 
ing a solution of the normal carbonate, which is used 
again for absorbing fresh quantities of the dioxide. 

Carbon Monoxide or Carbonic Oxide Gas: Formula — 
CO. Molecular Weight^-27.8. 
"When carbon burns with a limited supply of oxygen, 
carbonic oxide is formed. The production of this gas 
in an ordinary red-hot coal fire is often observed. Oxy- 
gen of the air, which enters at the bottom of the grate, 
combines with the carbon of the coal, forming carbon 
dioxide, this substance then passing upwards over the 
red-hot coals, parts with half its oxygen to the red-hot 
carbon, thus 

C0 2 +C=2CO. 

This carbon monoxide on coming out at the top of 
the fire meets with atmospheric oxygen, with which it 
at once combines, burning with a lambent blue flame, 
and re-forming carbon dioxide. Carbon monoxide can 
be prepared in various ways. It is formed when zinc 
oxide, ferric oxide, manganese dioxide, and many other 
oxides are heated with charcoal, it is also formed when 
chalk (calcium carbonate), magnesite (magnesium car- 
bonate), and other carbonates are heated with metallic 
zinc or iron filings, the decomposition which takes place 
in these cases is represented by the following equations, 
ZnO+C=Zn+CO. 
CaC0 3 +Zn=CaO+ZnO+CO. 

Carbon monoxide is also formed when carbon is 



CAEBON 95 

burnt in a limited quantity of oxygen, or by the action 
of carbon dioxide on heated charcoal. 
C0 2 +C=2CO. 

The latter reaction does not take place below 600°, 
and then only in presence of moisture. As the tem- 
perature is increased above 600°, moist carbon dioxide 
is partially converted into the monoxide, an equilibrium 
between the two gases being obtained at any particular 
temperature. 

When oxalic acid or an oxalate is heated with con- 
centrated sulphuric acid, a mixture of equal volumes 
of carbon monoxide and carbon dioxide is evolved, 
thus 

C 2 4 H 2 =H 2 0+CO+C0 2 . 

The carbon dioxide may readily be separated from 
the carbon monoxide either by passing the mixed gases 
through a solution of caustic soda or by collecting 
the mixture over water rendered alkaline by this sub- 
stance. 

Carbon and Sulphur Compounds. 
Carbon Bisulphide: Formula — CS^ Molecular Weight 

—75.57. 

Carbon bisulphide is prepared on a large scale by 
passing the vapor of sulphur over red-hot charcoal. 
For this purpose a large upright cast iron cylinder is 
employed, 10 or 12 feet long and one to two feet in 
diameter as in Fig. 8. This cylinder is placed above 
a furnace and surrounded by brickwork, and at the 
same time it is provided with a lid to admit of the 
whole being filled with charcoal. A second opening 
a, furnished with a hopper, exists at the bottom of the 
cylinder, and this serves to bring the sulphur into the 



96 



ELEMENTARY CHEMISTRY 



apparatus. The sulphur evaporates, and in the state 
of vapor combines with the red-hot carbon, impure 
carbon bisulphide distilling over by the tube c, and 
collecting in the vessel d under water. The tubes e 
serve as condensers, to separate the vapor of carbon 
bisulphide from sulphuretted hydrogen formed during 
the reaction, owing to the presence of hydrogen in 
the charcoal, the sulphuretted hydrogen being ab- 
sorbed by passing over the layers of slaked lime con- 




tained in the purifier f. In actua± work the yield of 
bisulphide is about 20 per cent, below the theoretical 
amount. The crude substance invariably contains 
sulphur in solution, from which, however, it may be 
separated by distillation, but other sulphur com- 
pounds, which impart to the crude material a most 
offensive odor, are also contained in the distillate. In 
order to remove these impurities, different processes 
are in use. The substance was formerly purified by 



CARBON 97 

frequent re-distillation over oil or fat, by which means 
the disagreeably smelling compounds were held back. 
Another means of purification employed is that of 
shaking the liquid with mercury, and allowing it to 
remain for a long time in contact with corrosive sub- 
limate in the cold, and then distilling it off white wax. 

Pure carbon bisulphide is a colorless, mobile, strong- 
ly refracting liquid, possessing a sweetish smell not 
unlike that of ether or chloroform. When carbon 
combines with sulphur to form carbon bisulphide, heat 
is absorbed which is given out again when it is de- 
composed. Like many substances of this class it is de- 
composed into its elements when the vapor is sub- 
jected to the shock caused by the explosion of fulmin- 
ate of mercury, but at atmospheric pressure the de- 
composition is only local and is not transmitted 
through the whole mass of the vapor. 

Carbon bisulphide is very inflammable and burns 
in the air, forming carbon dioxide and sulphur diox- 
ide. The combination with oxygen takes place slowly 
below the temperature of ignition, a slight phosphor- 
escence being observed, and as usual in such cases the 
exact temperature of ignition varies, but it occurs 
immediately at temperatures above 260°. The vapor 
of carbon bisulphide when mixed with oxygen can be 
exploded by a spark even when the mixture is thor- 
oughly dried, differing in this respect from carbon 
monoxide. With excess of oxygen a mixture of car- 
bon dioxide, sulphur dioxide, and trioxide, and some- 
times free sulphur is produced, carbon monoxide, car- 
bonyl sulphide, and unaltered carbon bisulphide being 
also present in the explosion wave. With an insuffi- 
cient quantity of oxygen, the products are carbon 



98 ELEMENTARY CHEMISTRY 

dioxide and monoxide, sulphur dioxide, carbonyl sul- 
phide and unaltered carbon bisulphide, but in no case 
is free carbon formed, and it is therefore improbable 
that the vapor is dissociated into its elements before 
combining with oxygen. 

CHLORINE. 



Symbol— 01. Atomic Weight— 35.2. Density— 35.2. 

Chlorine does not occur free in nature, but can 
easily be prepared from its compounds. It is found 
combined with metals forming chlorides, of these 
sodium chloride, sea- or rock-salt, is the most common. 
To obtain chlorine from this, it must be heated with 
sulphuric acid and manganese dioxide. 

Sodium chloride, sulphuric acid, and manganese di- 
oxide give chlorine, sodium sulphate, manganese sul- 
phate, and water. 

If one part by weight of salt to one part of man- 
ganese dioxide be mixed with two parts of sulphuric 
acid and two of water, and the mixture brought into 
a large flask, the chlorine gas is given off regularly 
upon the application of a very slight heat. In order 
to obtain the gas pure, it may be passed through the 
water contained in a wash-bottle before it is collected 
for use. Chlorine is a green-yellow gas, possessing a 
most disagreeable and peculiar smell, which, when the 
gas is present in small traces only, resembles that of 
seaweed, but when present in large quantities acts as 
a violent irritant, producing inflammation of the mu- 
cous membrane, and even causing death when inhaled. 
Chlorine gas when submitted to a pressure of five at- 



CHLORINE 99 

mospheres at the ordinary temperature is condensed 
to a heavy yellow liquid, but it has not yet been solidi- 
fied. This gas cannot be collected over water or mer- 
cury, as it is soluble in the former, 1 volume of water 
dissolving 2.37 volumes of chlorine at 15°, and it com- 
bines directly with the latter, forming mercuric chlo- 
ride. It can, however, be easily collected by displace- 
ment, as it is nearly 2.5 times as heavy as air. If 
metals in a freely divided state are brought into chlo- 
rine gas, they take fire spontaneously, forming metallic 
chlorides. Powdered arsenic, antimony, or thin copper 
leaf, burn when thrown into the gas. 

The most remarkable property of chlorine is its 
power of combining with hydrogen to form hydro- 
chloric acid. When these two gases are brought to- 
gether in equal volumes, they combine with explosion 
on bringing a flame into contact with them, or on ex- 
posing the mixture to sunlight. Chlorine is even able 
to decompose water in the sunlight, combining with 
the hydrogen and liberating the oxygen. Several ex- 
periments illustrative of .this property of chlorine may 
be mentioned. If a burning candle be plunged into 
this gas, the taper continues to burn, but with a very 
smoky flame, the hydrogen alone of the wax entering 
into combination with the chlorine, whilst the carbon 
is given off as smoke or soot. The same effect is pro- 
duced when a paper moistened with turpentine, a 
compound of carbon and hydrogen, is held in a jar 
of chlorine gas, the hydrogen of the turpentine at 
once combines with the chlorine, forming hydrochloric 
acid, and the carbon is liberated, so much heat is given 
off by this action that the paper frequently takes fire. 

The well-known bleaching 1 action of chlorine also de- 



100 ELEMENTARY CHEMISTRY 

pends upon its power of combining with the hydrogen 
of water and liberating the oxygen. Dry chlorine gas 
does not bleach, we may enclose a piece of cotton 
cloth or paper colored by a vegetable substance, as 
madder or indigo, in a bottle of dry chlorine, and no 
change of color takes place, even after the lapse of 
many weeks. If a few drops of water are added, the 
coloring matter is immediately destroyed, and the cot- 
ton or paper is bleached. Here the chlorine combines 
with the hydrogen of the water, and the oxygen at the 
moment of its liberation combines with the vegetable 
coloring matters, forming compounds destitute of 
color. Ordinary free oxygen has not this power, not 
at least to any great extent, and it is a frequent ob- 
servation that bodies in this nascent state have more 
active properties than the same bodies when in the 
free state. This difference depends upon the fact that 
the molecules, or smallest particles of an element which 
can exist in the free state, do not consist of the in- 
dividual atoms, but of a group of atoms. The mole- 
cule of a compound body contains two or more dis- 
similar atoms, whilst that of an element contains simi- 
lar atoms. The moment an element is liberated from 
a compound, the single atoms unite together to form a 
molecule, and the elementary body makes its appear- 
ance in the free state. If substances are present on 
which the element can act chemically, they are decom- 
posed by the chemical attractions of the liberated 
atom, which are more active in that state than when 
united to form a molecule. 

Chlorine is unable to bleach mineral colors. The 
difference between printers' ink, colored by lampblack 
or carbon, and writing ink, a vegetable black, is well 



CHLORINE 101 

illustrated by placing a sheet of paper having char- 
acters written and printed upon it in a solution of 
chlorine in water. Chlorine gas is largely used for 
bleaching purposes in the cotton, linen, and paper 
manufactures. It is sometimes used in the form of a 
gas, but more usually in combination with calcium 
and oxygen, forming the article called chloride of 
lime, a mixture of calcium chloride CaCl 2 , and cal- 
cium hypochlorite, CaCl 2 2 , or bleaching powder. 
Chlorine is also largely employed as a disinfectant 
and deodorant, its action on organic putrefactive sub- 
stances being similar to that upon organic coloring 
matters. 

Chlorine is usually prepared when a high degree of 
purity is not required, by the action of black oxide of 
manganese (manganese dioxide) on strong hydro- 
chloric acid, thus 

Mn0 2 +4HCl=Cl 2 +MnCl 2 +2H 2 0. 

When these two substances are brought together, a 
dark greenish-brown solution is first obtained, and 
this on heating evolves chlorine gas, whilst manganous 
chloride MnCl 2 is formed. There is little doubt that 
the dark-colored solution contains a higher unstable 
manganese chloride, probably MnCl 3 or MnCl 4 , which 
has not yet been isolated, and which on heating de- 
composes into chlorine and manganous chloride. 

In preparing the gas by this method the oxide of 
manganese should be in the form of small lumps free 
from powder, and the hydrochloric acid poured on so 
as to about cover the solid, on gently warming, the 
gas is copiously evolved. 

It is often more convenient for laboratory uses to 
liberate the hydrochloric acid in the same vessel in 



102 



ELEMENTARY CHEMISTRY 



which it is acted upon by the manganese dioxide, and 
for this purpose to place a mixture of one part of the 
latter substance and one part of common salt in a 
large flask as in Fig. 9, containing a cold mixture of 
two parts of strong sulphuric acid and two of water, 
on very slightly warming the mixture, a regular evo- 
lution of gas takes place. 




Fig. 9. 

Chlorine is not inflammable, and does not directly 
combine with oxygen, it unites, however, with great 
energy with hydrogen, forming hydrochloric acid, 
HC1, and to this property it owes its peculiar and 
valuable bleaching power. It also combines with many 
metals, giving rise to a class of compounds termed 
the metallic chlorides. 

In each case of combination with chlorine a definite 



CHLORINE 



103 



quantity of heat is given out, whilst sometimes light 
is also emitted, so that the essential phenomena of 
combustion are observed. Thus if we plunge a jet 
from which a flame of hydrogen burns into a cylinder 
of chlorine gas as in Fig. 10, the hydrogen continues 
to burn, but instead of water being produced, hydro- 
chloric acid is formed by the combustion. In like 




Fig. 10. 

manner, if we bring a light to the mouth (held down- 
wards) of a cylinder of hydrogen as in Fig. 11, and 
then bring this over a jet from which chlorine gas 
is issuing, a flame of chlorine burning in hydrogen 
will be seen. 

If two equal sized cylinders, filled, one with chlorine 
and the other with hydrogen, are brought mouth to 



104 



ELEMENTARY CHEMISTRY 



mouth, the two glass plates closing them withdrawn, 
and the gases allowed to mix, apd if then a flame is 
brought near the mouths of the cylinders, the mixed 
gases combine with a peculiar noise, and dense fumes 




Fig. 11. 



of hydrochloric acid gas are seen. This experiment 
must, however, be made in a room partially darkened, 
or performed by gas or candle-light, as the two gases 
combine with explosion in sunlight or strong daylight. 






CHROMIUM 105 

CHROMIUM. 
Symbol — Or. Atomic Weight — 51.7. 

Chromium is a substance whose compounds do not 
occur very widely distributed, or in large quantities, 
but they are much employed in the arts as pigments, 
many of them possessing a fine bright color. The 
chief ore of this metal is Chrome Ironstone, a com- 
pound isomorphous with Magnetic Oxide of Iron, 
found in America, Sweden, and the Shetlands, a com- 
pound lead chromate is also found in some quantity. 
Pure chromium appears to be the most infusible of 
all the metals, as it cannot be melted at a temperature 
sufficient to fuse and volatilize platinum. It has been 
obtained by another process in the form of bright 
crystals. 

Chromium and Chlorine Compounds. 
Chromic Chloride: Formula — Cr 2 Cl 6 . 

Chromic Chloride, the anhydrous chloride, is ob- 
tained as a sublimate, in beautiful violet crystals, by 
passing a current of chlorine gas over a red-hot mix- 
ture of chromium sesquioxide and charcoal. These 
crystals do not dissolve easily in water, but are readily 
soluble if a trace of chromium dichloride is present. 
The most ready way of preparing a solution of chro- 
mic chloride is to boil a solution of chromic acid or a 
chromate with hydrochloric acid and alcohol, the red 
or yellow solution after a few minutes being changed 
to a deep greenish-blue color. A solution of chromic 
sulphate may be obtained in the same way, by sub- 
stituting sulphuric acid for hydrochloric acid. Chro- 
mium sulphate forms a series of alums with potassium 



106 ELEMENTARY CHEMISTRY 

and ammonium sulphates, which have a deep purple 
tint, and are isomorphous with common alum. 

Chromium and Oxygen Compounds. 
Chromium Monoxide, CrO, Chromium Sesquioxide, 
Cr 2 3 and Chromium Trioxide, CrO s . 

Chromium Monoxide : Formula — CrO. 

Chromium Monoxide is only known in the hydrated 
state, as both it and its compounds absorb oxygen with 
great avidity. The hydrate is prepared as a brown 
precipitate by adding potash to the solution of chro- 
mium dichloride. 

Chromium Sesquioxide: Formula — Cr 2 3 . 

Chromium Sesquioxide, or Chromic Oxide is a dark 
green, perfectly stable powder, obtained by igniting 
the hydroxide formed by precipitating any soluble 
chromic salt with ammonia. It is employed as a green 
color for painting on porcelain, and produces the green 
of the emerald. A splendid green color is also ob- 
tained by heating potassium bichromate with boron 
trioxide, on dissolving in water a grass-green hydrox- 
ide remains behind, which is termed Guignet's green. 

Chromium Trioxide: Formula — Cr0 3 . 

Chromium Trioxide is obtained in the form of long 
ruby-red needle-shaped crystals by adding an excess 
of strong sulphuric acid to a concentrated solution of 
the bichromate. The crystals are very soluble in 
water, forming an acid solution of chromic acid. The 
excess of sulphuric acid may be removed by washing 
with concentrated nitric acid, and the crystals then 



COBALT 107 

dried in a current of air in a glass tube. The crystals 
of chromium trioxide are very easily reduced to ses- 
quioxide in presence of organic matter. So energetic 
is this evolution of oxygen, that ignition occurs when 
alcohol is dropped on the dry crystals. 

If a solution of chromium trioxide or of potassium 
bichromate is heated with hydrochloric acid, chromic 
chloride is formed and chlorine liberated, whereas, if 
chromium trioxide is heated with sulphuric acid, a 
chromic sulphate is formed and oxygen gas is given 
off. 

Chromic Acid and the * Chromates. 

If any chromic compound be fused with potassium 
carbonate, it becomes oxidized, and a soluble yellow 
chromate is formed, K 2 Cr0 4 . This is the mode in 
which the chromium compounds are prepared from 
chrome-iron ore. This yellow chromate is isomorphous 
with potassium, sulphate and manganate. When sul- 
phuric acid is added to a solution of this yellow salt 
in sufficient quantity to combine with half the base, 
large red crystals of the bichromate, K 2 Cr 2 7 , sep- 
arate out. This salt is largely used for the preparation 
of the chrome pigments. If to the solution of the bi- 
chromate a solution of chromium trioxide be added, a 
third salt, termed ter-chromate, K 2 Cr 3 O 10 , crystallizes 
out. 

COBALT. 
Symbol — Co. Atomic Weight — 59.0. 

Cobalt is a reddish-white, very tenacious metal, 
which is as infusible as iron, and, like the latter metal, 
is strongly magnetic. It is not found native, but oc- 
curs in combination with arsenic and sulphur, as two 



108 ELEMENTARY CHEMISTEY 

distinct minerals. The metal dissolves slowly in sul- 
phuric and hydrochloric acids with evolution of hy- 
drogen. The cobalt compounds are distinguished for 
the brilliancy of their color, they are employed as pig- 
ments, and they impart a magnificent blue tint to 
glass. 

The nitrate and sulphate of cobalt are also soluble 
salts, the latter is isomorphous with magnesium sul- 
phate. Cobalt sulphide is a black powder, insoluble 
in dilute acids. Cobalt compounds can be easily rec- 
ognized by the deep blue tint which very minute traces 
impart to glass, or to a borax bead, made by fusing 
borax into a colorless mass on the loop of a platinum 
wire. 

Cobalt and Chlorine Compounds. 
Cobalt Chloride: Formula — CoCl 2 . 

Cobalt chloride is a soluble salt obtained by acting 
on the oxide or on the metallic ore with hydrochloric 
acid, the solution yields on evaporation pink crystals 
of # the hydrated chloride, or, if further heated, blue 
crystals of the anhydrous salt. 

Cobalt and Oxygen Compounds. 

There are three oxides of cobalt, the monoxide, CoO, 
the sesquioxide, Co 2 3 and an oxide Co 3 4 , the former, 
on solution in acids, forms the series of cobalt salts, 
which are pink when hydrated, and blue when anhy- 
drous, while the sesquioxide does not form any salts. 
Cobalt monoxide, CoO, is obtained as a brown powder 
by carefully heating the rose-colored hydrate, precipi- 
tated by potash in solutions of cobalt, and Cobalt 



COPPER 109 

sesquioxide, Co 2 3 , is prepared by adding a solution 
of bleaching powder to a soluble protosalt. . 

COPPER. 

Symbol— Cu. Atomic Weight — 63.13. 

Copper is an important metal, largely used in the 
arts, and has been' known from early times, as it oc- 
curs native in the metallic state, and is moreover 
easily reduced from its ores. Metallic copper is found 
in considerable quantity in America and other locali- 
ties, but the chief sources of copper are the following 
ores. A compound of copper, sulphur, and iron, known 
as copper pyrites, the cuprous sulphide, the carbonate 
or malachite, and the red or cuprous oxide. Pure 
metallic copper can be obtained by reducing the oxide 
in a current of hydrogen gas, or by the electrolytic 
decomposition of a salt of copper. The process for 
obtaining copper on a large scale from the carbonate 
or oxide is a very simple one, merely reducing these 
ores together with carbon and some silica in a furnace. 
The reduction of the metal is more difficult when the 
commoner ore, copper pyrites, is employed. In this 
case the ore is repeatedly roasted, in order partially 
to convert the cuprous sulphide into oxide, and the 
roasted ore melted in a furnace with the addition of 
sand or silicious slag, in this operation the cuprous 
oxide becomes converted into the corresponding sul- 
phide, whilst the iron oxidizes and unites with the 
silica to form a light and fusible slag. The impure 
cuprous sulphide fuses and sinks to the lower portion 
of the furnace, forming the mat or coarse metal, and 
by repeating this operation a pure cuprous sulphide or 
fine metal is obtained. In order to prepare the metal- 



110 ELEMENTARY CHEMISTRY 

lie copper free from sulphur, this fine metal is roasted, 
and afterwards fused in contact with the air. During 
the first part of the operation a portion of the sulphur 
is burnt off, cupric oxide being formed, and in the 
later stages of the process this oxide acts upon the re- 
maining quantity of sulphide, forming sulphur diox- 
ide and metallic copper. In order to get rid of the 
last traces of oxide, the molten copper is poled or 
stirred up with a piece of green wood. 

Metallic copper possesses a peculiar deep red color, 
which is best seen when a ray of light is several times 
reflected from a bright surface of the metal, it is very 
malleable and ductile, and possesses great tenacity, it 
melts at a red heat, and is lightly volatile at a white 
heat, communicating a green tint to a flame of hydro- 
gen gas, and it is one of the best conductors of heat 
and electricity. Copper does not oxidize either in 
pure dry or moist air at ordinary temperatures, but 
if heated to redness in the air, it rapidly oxidizes to 
scales of copper oxide. Steam is not decomposed by 
metallic copper at a red heat. Finely divided copper 
dissolves in hydrochloric acid with evolution of hydro- 
gen, when heated with strong sulphuric acid, sulphur 
dioxide is evolved, and copper sulphate formed, and 
when acted upon with nitric acid, copper nitrate is 
produced, and nitric oxide liberated. 

Many of the copper alloys are of importance. Brass 
is an alloy containing about two-thirds of copper and 
one-third of zinc, it is harder than copper, and can 
be more easily worked. The addition of one to two 
per cent of lead improves the quality of brass for 
most purposes. Bronze, gun-metal, bell-metal, and 
speculum-metal are other alloys of copper and tin in 



COPPER 111 

varying quantities. They are all remarkable for the 
property of being hard and brittle when slowly cooled, 
but of becoming soft and malleable if they are cooled 
suddenly when red-hot by dipping into cold water. 

Copper and Chlorine Compounds. 
Chloride of Copper: Formula — CuCl 2 . 

Copper chloride is formed when copper is brought 
into chlorine gas, or when copper oxide is dissolved in 
hydrochloric acid; it forms green needle-shaped 
crystals, soluble in water and alcohol. The alcoholic 
solution burns with a characteristic green Same. 

Copper and Oxygen Compounds. 

Red Oxide, Cu 2 and Black Oxide or Copper Monox- 
ide, CuO. 

Cuprous or Red Oxide of Copper: Formula — Cu 2 0. 

This oxide occurs native in ruby-red octahedral crys- 
tals. It is artificially prepared by heating equivalent 
quantities of cupric oxide and copper filings, or by 
boiling a solution of copper sulphate and sugar, to 
which excess of caustic potash has been added, the 
sugar reduces the copper salt, and cuprous oxide is 
precipitated as a bright red powder. Cuprous oxide 
imparts to glass a splendid ruby-red color, it forms 
colorless salts with acids, which rapidly absorb oxygen 
from the air, and pass into the corresponding cupric 
compounds. The most important of these salts is cu- 
prous chloride, Cu 2 Cl 2 , a white solid substance obtained 
by dissolving a mixture of cupric oxide and metallic 
copper in hydrochloric acid. The solution of cuprous 



112 ELEMENTARY CHEMISTRY 

chloride possesses the remarkable property of absorb- 
ing carbonic oxide gas. 

Black Oxide of Copper: Formula — CuO. 

This oxide is formed when copper is heated in the 
air, or when copper nitrate is heated to redness, it 
yields the blue and green cupric salts, and it is largely 
used in the laboratory as a means of giving oxygen 
for the combustion of organic substances. Hydrated 
copper oxide is obtained as a light blue precipitate 
when caustic alkali is added to a cupric salt, when this 
is heated to 100°, it loses its water, and the anhydrous 
oxide falls as a brown powder. Cupric oxide is solu- 
ble in acids, furnishing a series of well crystallizing 
salts. 

Copper, Sulphur and Oxygen Compounds. 
Sulphate of Copper: Formula — CuS0 4 . 

This salt is sometimes known as blue vitriol, and is 
largely manufactured by dissolving copper oxide in 
sulphuric acid. It crystallizes in large blue crystals. 
When heated to redness, it loses all its water of crys- 
tallization, and forms a white powder, which again 
at a higher temperature decomposes, leaving copper 
oxide. Copper sulphate is employed in calico-printing, 
and in the manufacture of Scheele's and Brunswick 
green, and other copper pigments. 

Copper and Sulphur Compounds. 
Sulphide of Copper: Formula — CuS. ■ 

The insoluble copper salt, Copper sulphide, is ob- 
tained as a black precipitate, when sulphuretted hy- 



FLUORINE 113 

drogen gas is passed through an acidified solution of 
a copper salt. 

FLUORINE. 
Symbol— F. Atomic Weight— 18.9. 

Fluorine is a light greenish-yellow gas, paler and 
more purely yellow in color than chlorine, possessing 
a penetrating odor resembling that of hypochlorous 
acid. 

This element occurs combined with the metal cal- 
cium, forming calcium fluoride, or fluorspar, a mineral 
crystallizing in cubes. It also exists in large quanti- 
ties in cryolite, a mineral found in Greenland, whilst 
it has been detected in minute quantities in the teeth, 
and even in the blood, of animals. Fluorine is remark- 
able as forming no compounds with oxygen, and as 
being extremely difficult to prepare in a pure state. 
Many attempts have been made to obtain fluorine in 
the free state, but none of the methods which yield 
chlorine, bromine, or iodine give any result. By the 
action, however, of dry iodine upon dry silver fluoride, 
it appears that fluorine has been isolated, and it is 
found to be a colorless gas which does not act upon 
glass, and is absorbed by caustic potash with the for- 
mation of potassium fluoride and hydrogen dioxide. 

Fluorine is the most active element with which we 
are acquainted. It combines explosively with hydro- 
gen in the dark, the direct combination may be shown 
by simply inverting a jar filled with hydrogen over 
the positive exit tube of an electrolytic apparatus. As 
soon as the fluorine comes in contact with the hydro- 
gen a blue, red-bordered flame appears at the end of 
the platinum tube, hydrofluoric acid being formed 



114 ELEMENTARY CHEMISTRY 

which slowly attacks the glass jar. It decomposes 
water with the utmost avidity, uniting with the hydro- 
gen to form hydrofluoric acid, and liberating ozone, 
and therefore in all experiments with the gas the 
presence of moisture must be excluded as completely 
as possible. 

GOLD. 

Symbol— Au. Atomic Weight— 195.7. 

Gold is always found in the metallic state, it occurs 
in veins in the older sedimentary or in the plutonic 
rocks, and in the detritus of such rocks, it occurs in 
traces in the sand of most rivers, and although found 
generally in small quantities, it is a widely diffused 
metal. Previous to the discoveries of the gold-fields of 
California and Australia, it was obtained from certain 
iron pyrites. In order to obtain the gold, the detritus 
or sand which contains the metal is washed in a cradle 
or other arrangement, by means of which the lighter 
particles of mud or mineral are washed away, whilst 
the heavier grains of gold sink to the bottom of the 
vessel. When gold has to be worked in the solid rock, 
the mineral is crushed to powder, and then shaken up 
with mercury, and the gold thus extracted by amal- 
gamation. 

Gold possesses a brilliant yellow color, and, in thin 
films, transmits green light, it is nearly as soft as lead. 
It can be drawn out into fine wire, and is the most 
malleable of all the metals. It does not tarnish at any 
temperature, in dry or moist air, nor is it affected by 
sulphur, like silver, it is not acted upon by any single 
acid except selenic, but dissolves in presence of free 
chlorine and in nitro-hydrpchloric acid. At high tern- 



HYDROGEN 115 

perature gold is slightly volatile. Pure gold is best 
prepared by dissolving the ordinary metal in aqua 
regia, and adding ferrous sulphate, which is oxidized 
to ferric salt and precipitates the gold as a brown 
powder. 

Gold and Oxygen Compounds. 
Gold suboxide: Formula — Au 2 0. 
Gold trioxide: Formula — Au 2 3 . 

Gold unites with oxygen in two proportions, form- 
ing Gold suboxide, and Gold trioxide. Neither of these 
oxides forms salts with acids, but the latter unites 
with bases to form compounds called aurates, thus po- 
tassium aurate is KAu0 2 . Gold trioxide is obtained 
by adding zinc oxide or magnesia to a solution of gold 
chloride, the oxide falls as a brown powder, from 
which the zinc can be separated by nitric acid. Gold 
trioxide decomposes, in direct sunlight, into metal and 
oxygen, and is also reduced when heated to a tempera- 
ture of about 250°. The most important compound 
of gold trioxide is fulminating gold. This substance is 
obtained by acting on a solution of gold with excess 
of ammonia, a yellow-brown powder is precipitated, 
which, when dry, explodes very easily when heated to 
100°, or when struck with a hammer. 

HYDROGEN. 

Symbol — H. Atomic Weight — 1. Density — 1. 

Hydrogen is a colorless invisible gas, possessing 
neither taste nor smell, it is the lightest body known, 
being 14.47 times lighter than air. It occurs free in 
small proportions in certain volcanic gases, and it has 
been lately shown to exist absorbed in certain speci- 



116 ELEMENTARY CHEMISTRY 

mens of meteoric iron, but it is found in much larger 
quantities, combined with oxygen to form water, and 
it is by the decomposition of water, or of some other 
similar hydrogen compound, that the gas is always pre- 
pared. One-ninth of the weight of water consists of 
hydrogen, and this gas can readily be obtained from 
it by the action of certain metals, which decompose 
the water, combining with the oxygen to form a metal- 
lic oxide, and liberating the hydrogen as a gas. The 
metals of the alkalies, potassium and sodium, decom- 
pose water at the ordinary temperature of the air, 
some other metals, as iron, are only able to do so at a 
red heat, whilst others, for instance silver and gold, 
are unable to decompose water at all. When a small, 
piece of potassium is thrown into water, an instan- 
taneous decomposition of the water ensues, potassium 
hydroxide (caustic potash) is formed, and the hydro- 
gen of the water is liberated, so much heat being at 
the same time evolved that the hydrogen takes fire 
and burns. If the potassium, or, still better, sodium, 
be wrapped in a piece of wire gauze, as shown in Fig. 
12, and thus held in the water of the pneumatic trough, 
under the mouth of a cylinder, the hydrogen gas thus 
liberated may be collected, and its properties exam- 
ined. "Water consists of 2 parts by weight of hydro- 
gen and approximately 16 parts by weight of oxygen, 
and its chemical symbol is therefore H 2 0. When po- 
tassium or sodium act upon water, half the hydrogen 
is liberated, the metal taking its place, or water and 
potassium yield potassium hydroxide and hydrogen. 
This shows that for every 1 part by weight of hydro- 
gen which is liberated, 39.1 parts by weight of potas- 
sium enter into combination. The hydroxide which is 



HYDROGEN 



117 



formed dissolves in the water, but its presence can 
easily be detected either by the peculiar caustic taste 
which the solution possesses, whence its name, caustic 
potash, or by its power of turning to a blue color a 
solution of litmus which has been reddened by an acid. 
To prepare hydrogen by the action of red-hot iron 
on water, a wrought-iron pipe, like a gun-barrel, filled 




Fig. 12. 



with iron turnings, must be heated in a furnace, and 
steam from a small flask or boiler passed over the 
red-hot metal through the tube ; hydrogen gas is given 
off, and oxide of iron left in the tube. The most con- 
venient process of preparing pure hydrogen in quan- 
tity depends upon a property possessed by those met- 
als, such as iron or zinc, which decompose water at 



118 



ELEMENTARY CHEMISTRY 



a red heat, namely, that these metals are able to evolve 
hydrogen from water at the ordinary temperature of 
the air if a dilute acid be present. For the purpose 
of thus obtaining hydrogen, a flask or bottle is pro- 
vided with a cork and tube as represented in Fig. 13, 
some zinc clippings are introduced, and a mixture of 




Fig. 13. 



one part of sulphuric acid (a compound of sulphur, 
oxygen, and hydrogen) and eight parts of water poured 
in through the tube funnel. After a few minutes a 
rapid effervescence commences, and the evolved gas 
is collected over water in bottles or cylinders as in the 
case of oxygen. Care must, however, be taken that all 
the air is expelled from the flask before the hydrogen 



HYDROGEN 119 

is collected, this is easily ascertained to be the case by 
filling a test tube with the gas, and trying whether the 
gas burns quietly when a lighted candle is brought to 
the mouth of the tube held downwards. 

If we concentrate by boiling the liquid remaining 
in the flask after the evolution of the hydrogen, we 
find that white crystals separate out when the liquid 
cools, these consist of zinc sulphate. A given weight 
of zinc with sulphuric acid and water, can always 
be made to produce a certain weight of hydrogen, and 
a certain weight of zinc sulphate will .always be 
formed. It is found by experiment that 2 parts by 
weight of hydrogen can be obtained by dissolving 65.2 
parts of zinc with the formation of 161.2 parts of zinc 
sulphate. This can be represented by the equation — 

H 2 S0 4 +Zn=ZnS0 4 +H 2 , 
which not only indicates that sulphuric acid and zinc 
yield zinc sulphate and hydrogen, but also informs us 
as to the weights of the respective substances taking 
part in the reaction. 

Hydrogen burns in the air when a light is brought 
to it with a very slightly luminous, although extremely 
hot flame, and in the process the hydrogen combines 
with the oxygen of the air, forming water. The pro- 
duction of water by the combustion of hydrogen in the 
air may easily be shown by bringing a bright dry glass 
over the flame of hydrogen issuing from a fine jet, the 
glass becomes at once dimmed owing to the condensa- 
tion of water in small drops upon the cold dry sur- 
face. A number of these drops can be collected, and, 
upon examination, they are found to consist of pure 
water. Hydrogen does not support the combustion of 
a candle, nor the life of an animal. If a burning taper 



120 



ELEMENTARY CHEMISTRY 



is pushed up into a cylinder of this gas, held with its 
mouth downwards, the hydrogen burns at the mouth 
of the jar, while the taper is extinguished, it can, how- 
ever, be relit by the flame at the mouth. Hydrogen 
can be poured from one vessel to another in the air, 
but as it is lighter than air it must be poured upwards. 
The specific gravity of hydrogen, when air is taken 
as the unit, is found to be 0.0693, but for several rea- 
sons it is found to be more convenient to take hydro- 
gen itself as the unit, and compare the weight of the 
same volumes of other gases with hydrogen instead of 
air. One litre of hydrogen gas at 0° Centigrade and 
760 mm. pressure weighs 0.08936 gram. Free hydro- 
gen, like oxygen, has never been obtained in the liquid 
or solid state. 



Hydrogen and Chlorine Compounds. 
Hydrochloric or Muriatic Acid, HC1. 

Hydrochloric Acid: Formula — HC1. Molecular 

Weight— 36.2. 
This substance, the only known compound of chlo- 
rine and hydrogen, is obtained when equal volumes 
of chlorine and hydrogen are mixed and exposed to 
the diffused light of day, the gases then combine, and 
form an unaltered volume of hydrochloric acid gas. 
If the light be strong, this combination takes place so 
rapidly that a violent explosion occurs, owing to the 
sudden disengagement of heat consequent upon this 
combination. The volume of hydrochloric acid formed 
is equal to that of the chlorine and hydrogen, one 
molecule of hydrogen and one molecule of chlorine 
give two molecules of hydrochloric acid. 



HYDROGEN 121 

Hydrochloric acid may, however, be more easily pre- 
pared by heating common salt (sodium chloride) and 
sulphuric acid in a flask. 

Sodium chloride and sulphuric acid give hydrochlo- 
ric acid and hydrogen sodium sulphate. 

Hydrochloric acid is a colorless gas, 1.269 times 
heavier than air, it fumes strongly in damp air, com- 
bining with the moisture, and has a strongly acid re- 
action. It is very soluble in water, one volume of this 
liquid at 15° dissolving 454 volumes of the gas. This 
solution is the ordinary hydrochloric or muriatic acid 
of commerce. Under a pressure of 40 atmospheres the 
gas forms a limpid liquid. The gas can be collected 
over mercury, and its solubility in water strikingly 
shown by allowing a few drops of water to ascend to 
the surface of the mercury in contact with the gas. 
A rapid rise of the mercury in the jar immediately oc- 
curs. It fumes strongly in the air, and, when heated 
in a retort, loses at first hydrochloric acid gas, but 
after a time an aqueous acid distils over at the ordi- 
nary atmospheric pressure, containing 20.22 per cent 
of HC1, and boiling constantly at 110°. If the distilla- 
tion be conducted under a diminished pressure, the 
acid boils constantly at a lower temperature, and at- 
tains a composition which is different for each boiling 
point, hence the constant acids thus obtained by boil- 
ing the solution of hydrochloric acid gas in water can- 
not be considered as definite compounds of HC1 and 
water. This fact holds good for many other aqueous 
solutions of acids, etc., that residues constantly boil- 
ing at the same temperature, and having constant com- 
positions, are obtained on distillation, the composition 



122 ELEMENTARY CHEMISTRY 

and boiling point varying, however, with the pressure 
under which the distillation has been conducted. 

Enormous quantities of hydrogen chloride, common- 
ly called muriatic acid, from muria, sea salt, are ob- 
tained as a by-product in the manufacture of sodium 
carbonate. The acid thus produced is, however, very 
impure, having a yellow color, and containing iron, 
arsenic, organic matter, and sulphuric acid in solution. 

The exact composition of hydrogen chloride is best 
determined by decomposing the aqueous acid in the 
dark by means of a current of electricity, and collect- 
ing the gases (hydrogen and chlorine) evolved in a 
long tube, after allowing the decomposition to go on 
for some time. If the tube thus filled be opened in the 
dark under a solution of potassium iodide, the solution 
will rise in the tube, the iodine being liberated, the 
chlorine combining with the potassium, until exactly 
half the tube is filled with liquid, the remaining gas is 
found to consist of hydrogen. If the mixture of elec- 
trolytic gases, which can with care be sealed up in a 
strong tube having very finely drawn out ends, be ex- 
posed to the action of daylight, or of a bright artificial 
light, such as that of burning magnesium wire, imme- 
diate combination of the two gases will ensue, and on 
opening one of the ends under water this liquid will 
completely fill the whole of the tube, showing that the 
component gases were present in exactly the propor- 
tion needed to form hydrochloric acid gas which dis- 
solves in the water. 

Aqua Regia. Certain metals, such as gold and plati- 
num, and many metallic compounds, such as certain 
sulphides, which do not dissolve in either nitric or hy- 
drochloric acid separately, are readily soluble in a 



HYDROGEN 123 

mixture of both of these acids, especially upon warm- 
ing. This mixture has been termed Aqua Regia, and 
its solvent action depends upon the fact that it con- 
tains free chlorine, liberated by the oxidizing action 
of nitric acid on the hydrogen of the hydrochloric 
acid. The metals combine directly with this free chlo- 
rine to form soluble chlorides, and the sulphides are 
decomposed by it. The nitric acid is reduced to nitro- 
gen dioxide, and this combines with a portion of the 
chlorine to form the compounds NOC1 and NOCl 2 , 
which are liberated as yellowish gases, condensing to 
a dark yellow, very volatile liquid when they are led 
into a freezing mixture. The same compounds are 
formed by direct combination when two gases, nitro- 
gen dioxide and chlorine, are mixed together. If 
chlorine is present in excess, NOCl 2 is formed, while 
NOC1 is produced when the oxide of nitrogen is pres- 
ent in largest quantity. 



Hydrogen and Fluorine Compounds. 

Hydrofluoric Acid. 

Symbol— HP. Molecular Weight— 19.9. 

Hydrofluoric acid gas must be prepared in a leaden 
or platinum vessel, as glass is rapidly attacked by the 
vapor. The colorless gas thus obtained fumes strong- 
ly m the air, if it be passed into a metallic tube placed 
in a freezing mixture at the temperature of — 20°, a 
liquid is formed, this liquid is strong aqueous hydro- 
fluoric acid. It appears doubtful whether the acid has 
been obtained in the liquid state. The strong acid 
acts very violently upon the skin, producing painful 



124 ELEMENTARY CHEMISTRY 

wounds, and the fumes of the gas are likewise danger- 
ous from their corrosive power. When brought into 
contact with water the strong acid dissolves with a 
hissing noise. This aqueous acid attains a constant 
boiling point under the ordinary atmospheric pressure, 
when the liquid contains 37 per cent. 

The most remarkable property of hydrofluoric acid 
is its power of etching upon glass. This arises from 
the fact that fluorine forms, with the silicon contained 
in the glass. 

The acid thus obtained is a highly dangerous sub- 
stance, and requires the most extreme care in its manip- 
ulation, the inhalation of its vapor having produced 
fatal effects. A drop on the skin gives rise to blisters 
and sores which only heal after a very long period. 
From its great volatility the anhydrous acid can only 
be safely preserved in platinum bottles having a flanged 
mouth, a platinum plate coated with paraffin being 
tightly secured to the flanged mouth by clamp screws. 
The acid must be kept in a cool place not above a tem- 
perature of 15°, otherwise it is very likely to burst 
the bottle, and a freezing mixture should always be at 
hand when experimenting with it. 

Anhydrous hydrofluoric acid can also be obtained 
by acting on dry silver fluoride with hydrogen. 

If the hydrofluoric acid is not required to be 
perfectly anhydrous a much easier process than the 
foregoing can be adopted. This consists in the decom- 
position of fluor-spar by strong sulphuric acid, when 
hydrofluoric acid and calcium sulphate are formed. 

For this preparation vessels of platinum, or, on a 
large scale, vessels of lead, can be employed. 



HYDROGEN 



125 



Hydrogen and Oxygen Compounds. 

Water or Hydrogen Monoxide, H 2 and Peroxide of 

Hydrogen or Hydrogen Dioxide, H 2 2 . 

Water: Formula— H 2 0. Molecular Weight— 17.88. 
Density— 8.95. 
When hydrogen is burnt in the air, water is formed 
by the union of the hydrogen gas with the oxygen of 
the air. Cavendish first ascertained in 1871, that by 
the combustion of two volumes of hydrogen and one 




Fig. 14. 

of oxygen, pure water is formed and nothing else. 
Lavoisier in 1783, the French chemist, repeated and 
confirmed the experiments of Cavendish. The ap- 
paratus used by him for proving that hydrogen gas is 
really contained in water, is seen in facsimile in Fig. 
14. The water contained in the vessel a was allowed 
to drop slowly into the tube, e d, from which it flowed 
into the gunbarrel, d f, heated to redness in the fur- 



126 ELEMENTARY CHEMISTRY 

nace. Here part of the water was decomposed, the 
oxygen entering into combination with the metallic 
iron, whilst the hydrogen and some undecomposed 
steam passed through the worm, s, where the steam 
was condensed and the hydrogen was collected and 
measured in the glass bell-jar, m. The result of these 
experiments was found to be that 13.13 parts by weight 
of hydrogen united to 86.87 parts by weight of oxygen, 
or 12 volumes of oxygen with 22.9 volumes of hydro- 
gen. 

Bunsen's modification of Cavendish's method con- 
sisted in bringing known volumes of the two gases 
successively into a eudiometer, and combining them 
by means of an electric spark, carefully observing the 
consequent change in volume. The eudiometer is a 
strong glass tube c, Fig. 15, one metre in length and 
2.5 centimeters in diameter, closed at the top, open at 
the bottom, and having platinum wires sealed through 
the glass near the closed end. The tube is accurately 
graduated with a millimeter scale. The eudiometer 
contains at the top, one drop of water to render the 
gases moist. It is first completely filled with mer- 
cury and then inverted in the trough d, containing the 
same metal. 

Hj^drogen gas is now allowed to enter the tube, and 
the volume admitted measured, equal to 100 volumes, 
oxygen gas is next admitted, and the volume of the 
two mixed gases measured, 75 volumes of oxygen are 
added. In making this experiment, care must, how- 
ever, be taken that the temperature and atmospheric 
pressure are carefully measured by means of the ther- 
mometer b and the barometer a, shown in the figure, 
it is also necessary that the tube be not more than half 



HYDROGEN 



127 



full of the gaseous mixture, as great heat is evolved 
by the combustion, and hence a sudden expansion of 




Fig. 15. 



volume occurs, for which reason it is necessary to press 
down the open end of the tube upon a plate of soft 
rubber placed under the mercury. An electric spark 



128 ELEMENTARY CHEMISTRY 

is now passed through the gas along the platinum 
wires, when a flame is seen to pass down through the 
gas, showing that combination has occurred, the water 
produced will be deposited as dew upon the inside of 
the tube, and will then only take up about ^oVo P ar ^ 
of the bulk which its constituent gases occupied, so 
that its volume may be neglected. When the bottom 
of the eudiometer is opened, the column of mercury 
in the tube rises, and then only 25 volumes of gas re- 
main, and this turns out to be pure oxygen. Thus 100 
volumes of hydrogen require exactly 50 volumes of 
oxygen for their complete combustion. By a modifi- 
cation of this experiment, it can be shown that the 
volume of the gaseous water formed occupies exactly 
100 volumes, or 2 volumes of hydrogen unite with 1 
of oxygen to form 2 volumes of steam, hence the den- 
sity of steam or weight of 1 volume is — ^—=9. The 

A 

most striking method of demonstrating the composi- 
tion of water analytically is by splitting it up into its 
constituent gases by means of a current of electricity. 
For this purpose fill a glass vessel as in Fig. 16, with 
water acidulated with sulphuric acid to enable it to 
conduct the electricity, and bring two test tubes filled 
with water and inverted into this vessel over two 
small platinum plates attached to wires of the same 
metal passing through the soft rubber stopper at the 
bottom of the glass bowl, on connecting these with 
the terminals of a battery of three or four elements, 
an evolution of gas from each plate is noticed, that dis- 
engaged from the plate in connection with the carbon 
pole of the battery is found to be pure oxygen, whilst 
that coming off from the other plate connected with 



HYDROGEN 



129 



the zinc pole of the battery, is pure hydrogen gas. If 
the two tubes be graduated, it will be seen that the 
volume of the hydrogen is a very little more than 




Fig. 16. 

double that of the oxygen, owing to the oxygen being 
rather more soluble in water than the hydrogen. 

The water formed by the combination of the two 
gases can again be decomposed into free hydrogen and 
oxygen, and this can readily be effected by a current 



130 ELEMENTARY CHEMISTRY 

of electricity. For the purpose of exhibiting this pass 
a current of electricity from four or six elements by 
means of two platinum poles through some water acid- 
ulated with sulphuric acid as in Fig. 16. The instant 
contact is made, bubbles of gas begin to ascend from 
each platinum plate and collect in the graduated tubes, 
which at first are filled with the acidulated water. 
After a little time it will be seen that the plate which 
is in connection with the zinc of the battery evolves 
more gas than the one which is in contact with the 
carbon of the battery, and after the evolution has con- 
tinued for a few minutes one tube will be seen to con- 
tain twice as much gas as the other. On examination, 
the larger volume of gas will be found to be hydrogen, 
and will take fire and burn when a light is brought to 
the end of the tube in which it was collected, whilst 
the smaller volume of gas is seen to be oxygen, a glow- 
ing chip of wood being rekindled when plunged into 
it. 

Whereas the combination of hydrogen with oxygen 
is attended by the evolution of energy in the forms of 
light and heat, the inverse change, the decomposition 
of water, requires energy to be supplied, and in the 
experiment just described, this is conveyed to the water 
in the form of the energy of the electrical current. 

In order to collect the mixed explosive gases evolved 
by the electrolytic decomposition of water, an ap- 
paratus such as is shown in Fig. 17 is used. 

Oxygen being 16 (exactly 15.88, Hydrogen as 1) 
times as heavy as hydrogen, and these gases combin- 
ing to form w r ater in the proportions by volume of one 
volume of the former to two of the latter, we now 
know that the proportions by weight in which these 



HYDROGEN 



131 



gases exist in water must be as 16 to 2. It is never- 
theless most important that this calculation be verified 
by direct experiment. For this purpose, use is made 
of the fact that copper oxide when heated alone does 
not part with any of its oxygen, but when heated in 




Fig. 17, 



presence of hydrogen it parts with as much oxygen 
as will, by combining with the hydrogen, form water, 
being itself wholly or partly reduced to metallic cop- 
per. If, therefore, a known weight of copper oxide 



132 ELEMENTARY CHEMISTRY 

is taken, heat and pure hydrogen passed over it until 
it has parted with all its oxygen, and if we collect 
and weigh all the water thus formed, and likewise 
weigh the remaining metallic copper, we shall have 
made a synthesis by weight of water. For the loss in 
weight of the copper oxide is the weight of oxygen 
which has combined with hydrogen to form water, and 
the difference between this weight and that of the 
water formed, is the weight of the hydrogen thus com- 
bined. The arrangement used for this determination 
is represented in Fig. 18. The hydrogen evolved by 
zinc from sulphuric acid in the bottle on the left hand 
is purified from any trace of arsenic, sulphur, and 
moisture which it may contain by passing through the 
U tubes, containing absorbent substances. The tube 
containing a very hygroscopic substance, is weighed 
both before and after the experiment, and if no in- 
crease occurs, the dryness of the gas is insured. The 
gas then comes in a perfectly pure state into contact 
with the heated copper oxide contained in the bulb B. 
This first bulb, which is accurately weighed, is placed 
in connection with a second bulb B, in which the water 
formed by the reduction of the oxide collects, any 
moisture which may escape condensation in this bulb 
is retained in weighed drying tubes, containing frag- 
ments of pumice stone moistened with sulphuric acid. 
Most careful experiments made according to this meth- 
od, carried out with many precautions which cannot 
here be detailed, have shown that 88.89 parts of oxy- 
gen by weight unite with 11.11 parts of hydrogen to 
form 100 parts of water. 

Free oxygen and hydrogen combine together, when 
a light is brought in contact with them, with so much 



HYDROGEN 



133 




Fig. 18. 



134 ELEMENTARY CHEMISTRY 

force that a violent and dangerous explosion occurs 
from the sudden expansion caused by the great heat 
evolved in combination. If we fill a strong soda-water 
bottle one-third full with oxygen and two-thirds with 
hydrogen, and then bring a flame to the mouth, the 
gases combine, producing a sudden detonation like the 
report of a pistol. Many fatal accidents have occurred 
to persons who have carelessly experimented with 
large volumes of this explosive mixture. In order to 
exhibit the great heat evolved by the combination of 
the two gases, the oxyhydrogen blowpipe is employed, 
in this arrangement the gases are contained separate- 
ly in two rubber bags, being only brought together at 
the point at which the combination is desired, so that 
all danger of explosion is avoided. The flame thus pro- 
duced is very slightly luminous, but its temperature is 
so high, that the most difficultly fusible metals, such 
as platinum, may be easily melted in it, whilst iron 
wire held in the flame burns with beautiful scintilla- 
tions, forming an oxide of iron. A piece of chalk or 
lime placed in this flame becomes heated to bright 
whiteness and emits an intense light, often used for 
signal purposes. 

Water exists in nature in three forms : In the solid 
form as ice, in the liquid state as water, and in the 
gaseous form as steam. At all temperatures between 
0° and 100° Centigrade it takes the liquid form, and 
above 100° it entirely assumes the gaseous form, under 
the ordinary atmospheric pressure of 760 mm. The 
melting point of ice is always found to be a constant 
temperature, and hence it is taken as the zero of the 
Centigrade scale, water may, however, under certain 
conditions, be cooled below 0° Centigrade without be- 



HYDROGEN 135 

coming solid, still ice can never exist at a temperature 
above 0° Centigrade. In passing from the solid to the 
liquid state water becomes reduced in volume, and on 
freezing a sudden expansion, from 1 volume to 1.099 
takes place. That this expansion exerts an almost ir- 
resistible force is well illustrated by the splitting of 
rocks during the winter. Water penetrates into the 
cracks and crevices of the rocks, and on freezing 
widens these openings, this process being repeated over 
and over again, the rock is ultimately split into frag- 
ments. Hollow balls of thick castiron can thus easily 
be split in two by filling them with water and closing 
by a tightly-fitting screw, and then exposing them to 
a temperature below 0° Centigrade. 

In the passing from solid ice to liquid water, we not 
only observe this alteration in bulk, but we notice that 
a very remarkable absorption, or disappearance of 
heat, occurs. This is rendered plain by the following 
simple experiment : Take a kilogram of water at the 
temperature 0°, and another kilogram of water at 79°, 
if we mix these, the temperature of the mixture will 
be the mean, or 39°. 5, if, however, we take 1 kilogram 
of ice at 0° and mix it with a kilogram of water at 
79°, we shall find that the whole of the ice is melted, 
but that the temperature of the resulting 2 kilograms 
of water is exactly 0°, in other words, the whole of the 
heat contained in the hot water has just sufficed to 
melt the ice, but has not raised the temperature of the 
water thus produced. Hence we see that in passing 
from the solid to the liquid state a given weight of 
water takes up or renders latent just so much heat 
as would suffice to raise the temperature of the same 
weight of water through 79° Centigrade, the latent 



136 ELEMENTARY CHEMISTRY 

heat of water is therefore said to be 79 thermal units 
— a thermal unit meaning the amount of heat required 
to raise a unit weight of water through 1° Centigrade. 
When water freezes, or becomes solid, this amount of 
heat which is necessary to keep the water in the liquid 
form, and is therefore well termed the heat of liquidi- 
ty, is evolved, or rendered sensible. A similar disap- 
pearance of heat on passing from the solid to the liquid 
state, and a similar evolution of heat on passing from 
the liquid to the solid form, occurs with all substances, 
the amount of heat thus rendered latent or evolved 
varies, however, with the nature of the substance. A 
simple means of showing that heat is evolved on solidi- 
fication consists in obtaining a saturated hot solution 
of Glauber's salt, sodium sulphate, and allowing it to 
cool. Whilst it remains undisturbed, it retains the 
liquid form, but if agitated it at once begins to crys- 
tallize, and in a few moments becomes a solid piass. 
If a delicate thermometer be now plunged into the 
salt while solidifying, a sudden rise of temperature 
will be noticed. Similarly water at rest may be cooled 
down below 0° Centigrade without solidifying, but if 
agitated it at once solidifies, and the temperature of 
the whole mass instantly rises to 0° Centigrade. 

When water is heated from 0° to 4°, it is found to 
contract, thus forming a striking exception to the gen- 
eral law, that bodies expand when heated and con- 
tract on cooling, on cooling from 4° to 0° it expands 
again. Above 4°, however, it follows this ordinary 
law, expanding when heated, and contracting when 
cooled. This peculiarity in the expansion and contrac- 
tion of water may be expressed by saying that the 
point of maximum density of water is 4° Centigrade, 



HYDROGEN 137 

that is, a given bulk of water will at this temperature 
weigh more than at any other. Although the amount 
of contraction on heating from 0° to 4° is but small, 
1 volume of water at 4° becoming 1+0.00012 at 0°, it 
yet exerts a most important influence upon the econo- 
my of nature. In order better to understand what the 
state of things would be if water obeyed the ordinary 
laws of expansion by heat, we may perform the fol- 
lowing experiment: Take a jar containing water at 
a temperature above 4°, place a thermometer at the 
top and another at the bottom of the liquid. Now 
briLg the jar into a place where the temperature is 
below the freezing point, and observe the temperature 
at the top and bottom of the liquid as it cools. It will 
be seen that at first the upper thermometer always 
indicates a higher temperature than the lower one, 
after a short time both thermometers mark 4°, and, 
as the water cools still further, it will he seen that the 
thermometer at the top always indicaxes a lower tem- 
perature than that shown by the one at the bottom. 
This cooling goes on till the temperature of the top 
layer of water sinks to 0°, after which a crust of ice 
is formed, jut if the mass of the water be sufficiently 
large, the temperature of the water at the bottom is 
never reduced below 4°. In nature precisely the same 
phenomenon occurs in the freezing of lakes and rivers, 
the surface-wrter is gradually cooled by cold winds, 
and thus becoming heavier, sinks, whilst lighter and 
warmer water rises to supply its place, this goes on til] 
the temperature of the whole mass is reduced to 4°, 
after which the surface-water never sinks, however 
much it be cooled, as it is always lighter than the 
deeper water at 4°. Hence ice is formed only at the 



138 ELEMENTARY CHEMISTRY 

top, the mass of water retaining the temperature of 
4°. Had water become heavier as it cooled down to 
the freezing point, a continual circulation would be 
kept up until the whole mass was cooled to 0°, when 
solidification of the whole would ensue. Thus our 
lakes and rivers would be converted into solid masses 
of ice, which the summer's warmth would be quite in- 
sufficient thoroughly to melt; hence the climate of our 
now temperate zone might approach in severity that of 
the Arctic regions. Sea-water does not freeze en 
masse, owing to the great depth of the ocean, which 
prevents the whole from ever being cooled down to 
the freezing point, similarly very deep lakes never 
freeze, as the temperature of the whole mass never 
gets reduced to 4° Centigrade. 

In passing from the liquid to the gaseous state, water 
exhibits several interesting and important phenomena. 
In the first place, when we heat water to 100° C. 
it begins to boil, or enters into ebullition, that is, a 
rapid disengagement of water-gas, or steam, from the 
lower or most heated surface takes place, this is well 
seen when water is heated in a glass globe over a 
gas flame. In this passage from the liquid to the gaseous 
state, a large quantity of heat becomes latent, the tem- 
perature of the steam given off being the same as that 
of the boiling water, as, like all other bodies, water 
requires more heat for its existence as a gas than as 
a liquid. The amount of heat latent in steam is rough- 
ly ascertained by the following experiment. . Into 1 
kilogram of water at 0°, steam from boiling water, 
having the temperature of 100°, is passed until the 
water boils, it is then found that the whole weighs 
1.187 kilos., or 0.187 kilo, of water in the form of 



i 



HYDROGEN 139 

steam at 100° has raised 1 kilo, of water from 0° to 
100° ; or 1 kilo, of steam at 100° would raise 5.36 kilos, 
of ice-cold water through 100°, or 536 kilos, through 
1°. Hence the latent heat of steam is said to be 536 
thermal units. 

Whenever water evaporates or passes into the gase- 
ous state, heat is absorbed, and so much heat may be 
thus abstracted from water that it may be made to 
freeze by its own evaporation. A beautiful illustra- 
tion of this is found in an instrument consisting of a 
bent tube, having a bulb on each end, and containing 
water and vapor of water, but no air. On placing all 
the water in one bulb, and plunging the empty bulb 
into a freezing mixture, a condensation of the vapor 
of water in this empty bulb occurs, and a correspond- 
ing quantity of water evaporates from the other bulb 
to supply the place of the condensed vapor, this con- 
densation and evaporation go on so rapidly that in a 
short time the water cools down below 0°, and a solid 
mass of ice is left in the bulb. By a very ingenious 
arrangement this plan of freezing water by its own 
evaporation has been practically carried out on a large 
scale by M. Carre, by means of which ice can be most 
easily and cheaply prepared. This arrangement con- 
sists simply of a powerful air-pump A, Fig. 19, and a 
reservoir B of a hygroscopic substance, such as strong 
sulphuric acid. On placing a bottle of water in con- 
nection with this apparatus, and on pumping for a 
few minutes, the water begins to boil rapidly, and the 
temperature of the water is cooled so low by its own 
evaporation as to freeze to a mass of ice. 

Water, and even ice, constantly give off steam or 
aqueous vapor at all temperatures, when exposed to 



140 



ELEMENTARY CHEMISTRY 



the air, thus we know that if a glass of water be left 
in a room for some days, the whole of the water will 
gradually evaporate. This power of water to rise in 
vapor at all temperatures is called the elastic force, 
or pressure, of aqueous vapor, it may be measured, 
when a small quantity of water is placed above the 
mercury in a barometer, by the depression which the 




Fig. 19. 



tension of the vapor thus given off is capable of ex- 
erting upon the mercurial column. If we gradually 
heat the drops of water thus placed in the barometer, 
we shall notice that the column of mercury gradually 
sinks, and when the water is heated to the boiling 
point, the mercury in the barometer tube is found to 
stand at the same level as that in the trough, showing 
that the elastic force of the vapor at that temperature 



HYDROGEN 141 

is equal to the atmospheric pressure. Hence water 
boils when the pressure of its vapor is equal to the 
superincumbent atmospheric pressure. On the tops 
of mountains, where the atmospheric pressure is less 
than at the sea's level, water boils at a temperature 
below 100°. At Quito, where the mean height of the 
barometer is 527 mm., the boiling point of water is 
90°.l, that is, the pressure of aqueous vapor at 90°. 1 
is equal to the pressure exerted by a column of mer- 
cury 527 mm. high. Founded on this principle, an in- 
strument has been constructed for determining heights 
by noticing the temperatures at which water boils. A 
simple experiment to illustrate this fact consists in 
boiling water in a globular flask, into the neck of which 
a stopcock is fitted, as soon as the air is expelled, the 
stopcock is closed, and the flask removed from the 
source of heat, the boiling then ceases, but on immers- 
ing the flask in cold water, the ebullition recommences 
briskly, owing to the reduction of the pressure conse- 
quent upon the condensation of the steam, the pressure 
of the vapor at the temperature of the water in the 
flask being greater than the diminished pressure. All 
other liquids obey a similar law respecting ebullition, 
but as the pressure of their vapors are very different, 
their boiling points vary considerably. 

When steam is heated alone, it expands according 
to the law previously given for permanent gases, but 
when water is present, and the experiment is per- 
formed in a closed vessel, the elastic force of the steam 
increases in a far more rapid ratio than the increase of 
temperature. The following table gives the pressure 
of aqueous vapor, as determined by experiment, at 
different temperatures measured on the thermometer. 



142 



ELEMENTARY CHEMISTRY 



Pressure of the 


Vapor of Water. 








Pressure in atmos- 


Temperature 


Tension in milli- 


Temperature 


pheres, 1 atmos- 


Centigrade. 


metres of mercury. 


Centigrade. 


phere =760 mm. 
of mercury. 


-20° 


0.927 


100° 


1. 


-10 


2.093 


111.7 


1.5 





4.600 


120.6 


2. 


+5 


6.534 


127.8 


2.5 


10 


9.165 


133.9 


3. 


15 


12.699 


144.0 


4. 


20 


17.391 


159.2 


6. 


30 


31.548 


170.8 


8. 


40 


54.906 


' 180.3 


10. 


50 


91.982 


188.4 


12. 


60 


148.791 


195.5 


14. 


70 


233.093 


201.9 


16. 


80 


354.280 


207.7 


18. 


90 


525.450 


213.0 


20. 


100 


760.000 


224.7 


25. 



We now see why the barometric height must be no- 
ticed in graduating a thermometer, if the height differ 
from 760 mm. the temperature of the water boiling 
under that pressure will not be quite 100° Centigrade. 
A metal vessel is here employed, because it is found 
that water does not always boil at 100° in glass ves- 
sels, even though the atmospheric pressure be 760 mm., 
owing to some molecular action analogous to cohesion 
between the glass and water. 

Pure water and ice, when seen in large masses, are 
found to possess a blue color. In order to obtain pure 
water the chemist is obliged to distill river or spring- 
water, that is ; to boil the water and collect the water 
formed by the condensation of the steam thus pro- 
duced, as all such water contains more or less solid 



HYDROGEN, 



143 



matter in solution derived from the surface of the 
earth over which the water flows, this dissolved solid 
matter is left behind on boiling off the water. Solid 
matter in suspension can be got rid of by the simpler 
process of filtration through paper, sand, etc. An ar- 
rangement for distillation is seen in Fig. 20, it con- 
sists of a retort B in which the impure water is placed, 
connected with a condenser D, made of tubes, between 
which a current of cold water is made to circulate. 




Fig. 20. 



The distilled water is collected in the flask placed 
at the end of the apparatus. Rain-water is the purest 
form of water occurring in nature, but even this con- 
tains impurities derived from the dust, etc., in the air, 
and no sooner does it touch the earth's surface than 
it dissolves some of the materials with which it comes 
in contact, and according to the nature of the ground 
over which it passes becomes more or less impure. All 
fresh-water on the earth's surface has been derived 
from the ocean by a vast process of distillation, hav- 



144 ELEMENTARY CHEMISTRY 

ing been deposited in the form of rain or snow from 
the atmosphere. 

All the rain-water ultimately passes in the form of 
spring-water, or river-water, into the sea, carrying 
with it the soluble constituents which have been dis- 
solved out from the strata through which it has perco- 
lated. In consequence of this continual accession of 
soluble salts, and removal of pure water by evapora- 
tion, the sea-water is rendered salt, it contains about 
35 parts of solid matter, 28 of which consist of common 
salt or sodium chloride in solution in 1,000 parts of 
water. 

Water is the most general solvent for chemical sub- 
stances with which we are acquainted. Most salts are 
soluble to a greater or less extent in water, and are 
deposited again in crystals when the water is evapo- 
rated, we are unacquainted with any simple general 
law regulating the quantities of salts taken up by 
water, in most cases the solubility is greater in hot 
than in cold water. Water is also contained in the 
solid state in combination as water of crystallization 
in many salts, when this water is driven off by heat, 
the crystal falls to powder. Gases also dissolve in 
water in quantities varying with the nature of the 
gas, the temperature, and' the pressure to which the 
gas is subjected. It is solely in consequence of the 
presence of oxygen derived from the air dissolved in 
the water of lakes, rivers, and seas, that fish are en- 
abled to keep up their respiration, as the water passes 
through their gills the oxygen is taken up to purify 
their blood. 



HYDROGEN 145 



Peroxide of Hydrogen or Hydrogen Dioxide: For- 
mula— H r 2 . Weight— 33.76. 

This substance has received the name of oxygenated 
water, as it easily decomposes into oxygen and water. 
It is found to contain twice as much oxygen as water 
does, consisting of 2 parts by weight of hydrogen com- 
bined with 32 of oxygen, hence, if we represent water 
by the symbol H 2 0, hydrogen dioxide will be written 
H 2 2 . It does not occur in nature, but is artificially 
prepared by acting on barium dioxide, Ba0 2 , with hy- 
drochloric acid, H 2 C1 2 , an exchange takes place be- 
tween the barium and hydrogen, giving rise to hydro- 
gen dioxide and barium chloride. 

Hydrogen 'dioxide may also be prepared by passing 
carbonic acid gas through barium dioxide suspended 
in w r ater, when barium carbonate separates out as a 
white powder insoluble in water, and hydrogen dioxide 
remains in solution. The reaction is represented by 
the following equation : 

Ba0 2 +H 2 0+C0 2 =BaC0 3 +H 2 2 . 

The aqueous solution of the dioxide is concentrated 
by allowing the wrater to evaporate under the receiver 
of an air-pump, the liquid after a time becomes thick, 
but it cannot be entirely freed from water. Hydro- 
gen dioxide is chiefly characterized by the ease with 
which it loses half its oxygen, this gas is slowly given 
off at 20° ? but at 100° Centigrade the evolution of oxy- 
gen becomes very rapid. In consequence of the readi- 
ness with which it gives off oxygen, hydrogen dioxide 
acts as a powerful bleaching agent, rapidly oxidizing 
and destroying vegetable coloring matter. A remarka- 



146 ELEMENTARY CHEMISTRY 

hie decomposition occurs when this substance is 
brought in contact with ozone, common oxygen and 
water being produced. Another interesting reaction 
occurs when silver oxide is brought together with hy- 
drogen dioxide, as the silver oxide is reduced to metal- 
lic silver whilst water and common oxygen are formed. 

Hydrogen and Sulphur Compounds. 

Hydrogen Sulphide, H 2 S and Hydrogen Disulphide, 

H 2 S 2 . 

Hydrogen Sulphide or Sulphuretted Hydrogen: For- 
mula— H 2 S. Molecular Weight— 33.83. 

This gas is best prepared by the action of dilute sul- 
phuric acid upon iron sulphide, PeS, iron sulphate be- 
ing also formed, thus: 

FeS+H 2 S0 4 =FeS0 4 +H 2 S, 
where two atoms of hydrogen change place with one 
of divalent iron. Pig. 21 represents a convenient form 
of apparatus for the production and purification of 
this gas. It may be collected over warm water, and 
is a colorless gas, possessing the peculiar odor of rot- 
ten eggs, it burns on application of a light with a 
bluish flame, forming water and sulphur dioxide. 
When inhaled, it acts as a poison on the animal 
economy, even if diluted with large quantities of air. 
Sulphuretted hydrogen gas dissolves in water to a 
considerable extent, imparting its peculiar smell and 
a slightly acid reaction to the water. One volume of 
water at 0° dissolves 4.37 volumes of the gas, whilst at 
15° 3.23 volumes are soluble! Exposed to a tempera- 



HYDROGEN 



147 



ture of — 74°, this gas condenses to a colorless, mobile 
liquid, which, when further cooled to — 85° freezes 
to a transparent ice-like solid. Under a pressure of 
about seventeen atmospheres this gas liquefies at the 
ordinary temperature of the air. Sulphuretted hydro- 
gen occurs free in nature in volcanic gases, as well as 
in the water of certain springs. Harrogate waters owe 
their peculiar odor and medicinal power to the pres- 
ence of this gas. It is likewise generated by the 




Fig. 21. 

putrefaction of animal matters, such as albumen, or 
the white of eggs, which contains sulphur, also by the 
deoxidation of sulphates in presence of decaying or- 
ganic matter. 

The composition of sulphuretted hydrogen may be 
ascertained either by heating a small piece of metallic 
tin in a known volume of the gas, when tin sulphide 



148 ELEMENTARY CHEMISTRY 

will be formed, and hydrogen liberated, or by decom- 
posing the gas by means of a red-hot platinum wire, 
when the whole of the sulphur is deposited, and hy- 
drogen set free. In both cases the volume of hydro- 
gen obtained is found to be equal to that of the gas 
employed, and hence 2 volumes of sulphuretted hydro- 
gen, weighing 34, consist of 2 volumes of hydrogen, 
weighing 2, and 1 volume of sulphur vapor, weighing 
32. 

Sulphuretted hydrogen is an invaluable re-agent 
in the laboratory, as by its means we are enabled to 
separate the metals in groups. If we pass a current 
of this gas through a solution of a copper salt, to 
which a small quantity of acid has been added, we 
obtain an immediate precipitate of copper sulphide, 
thus 

CuS0 4 +H 2 S=CuS+H 2 S0 4 . 
If we do the same with a solution of an iron salt, we 
get no such precipitate, because iron sulphide is solu- 
ble in an acid, but on the addition of an alkali, iron 
sulphide is at once precipitated, thus 

FeS0 4 +2KHO+H 2 S=FeS+K 2 S0 4 +2H 2 0. 

We may thus divide the metals into groups: First, 
those which, like copper are precipitated by sulphur- 
etted hydrogen from an acid solution, or the copper 
group, second, those, which are not precipitated by 
sulphuretted hydrogen in an acid solution, but which 
are so precipitated in an alkaline one, or the iron 
group v and third, those which are in no case precipi- 
tated by this re-agent, as their sulphides are soluble 
either in water, acids, or alkalies, to this group be- 
long the metals of the alkalies and alkaline earths. 



HYDROGEN 149 



Hydrogen Disulphide: H 2 S 2 . 

This substance is an oily liquid obtained by pour- 
ing a solution of calcium disulphite into hydrochloric 
acid. This compound is not used in the arts, but it is 
interesting as having a composition corresponding to 
H o 2 , and because, like this substance, it possesses 
bleaching properties. It easily decomposes into sulphur- 
etted hydrogen and free sulphur. 

Hydrogen, Carbon and Nitrogen Compounds. 
Hydrocyanic or Prussic Acid: Formula — HON. 

Hydrocyanic acid is generally prepared from potas- 
sium ferrocyanide, which is decomposed by dilute sul- 
phuric acid. The process enables us to prepare any 
desired quantity, either of the aqueous acid of dif- 
ferent strengths, or of the anhydrous compound. A 
cold mixture of 14 parts of water and 7 parts of con- 
centrated sulphuric acid is poured upon 10 parts of 
coarsely powdered potassium ferrocyanide contained 
in a large retort the neck of which is placed upwards 
in a slanting direction. If the anhydrous acid be re- 
quired, the vapor is allowed to pass through cylinders 
or U-tubes filled with calcium chloride and surrounded 
by water heated to 30°, after which the gas, thus 
dried, is condensed by passing into a receiver sur- 
rounded by ice or a freezing mixture. To prepare the 
aqueous acid the drying apparatus may be omitted 
and the gas evolved from a flask as shown in Fig. 22. 
passed through a Liebig's condenser, and then led into 
the requisite quantity of distilled water. The action 



150 



ELEMENTARY CHEMISTRY 



of the dilute sulphuric acid upon potassium ferrocy- 
anide is represented by the following equation 

2Fe (CN) 6 K 4 +3H 2 S0 4 =3K 2 S0 4 +Fe (CN) 6 K 2 Fe+ 

6HCN. 

Pure hydrocyanic acid is one of the most powerful 
and rapid of known poisons. When a small quantity 
of the vapor of the pure substance is drawn into the 
lungs instant death ensues, small quantities produce 




Fig. 22. 



headache, giddiness, nausea, dyspnoea, and palpitation. 
A few drops brought into the eye of a dog kill it in 
thirty seconds, whilst an internal dose of 0.05 grain is 
usually sufficient to produce fatal effects upon the 
human subject, but cases have been known in which 
0.1 grain has been taken without death ensuing. As 
antidotes, ammonia and chlorine water have been 
proposed, and these appear to be efficacious, although 
we are unable to explain their mode of action, for 



HYDROGEN 151 

ammonia under ordinary conditions only forms am- 
monium cyanide, and chlorine cyanogen chloride, both 
of which are bodies as poisonous as prussic acid itself. 
Hydrocyanic acid is used as a medicine, and is a coii- 
lent of several preparations, such as laurel water 
and bitter-almond water which are obtained by dis- 
tilling the leaves of the common laurel, or bitter al- 
monds with water. 

Hydrogen, Boron and Oxygen Compounds. 
Boric or Boracic Acid: Formula — H BO.. 

Boric acid is manufactured from minerals, such as 
borocalcite. which occurs in considerable quantities in 
the nitre beds of Peru and Chili. It is likewise pre- 
pared from the natural borax or tineab which was 
first obtained from the basins of dried-up lagoons in 
Central Asia, and has been found in the borax lake in 
California in such quantities that the amount there 
obtained is sufficient to supply the whole demand of 
the United States. For the purpose of preparing boric 
acid from these sources, the minerals are dissolved in 
hot hydrochloric acid, the boric acid, which separates 
out on cooling, being recrystallized from hot water. 

Boric acid crystallizes from aqueous solution in 
shining six-sided lamina unctuous to the touch, and 
belonging to the triclinic system. 

Boric acid is a weak acid, and its cold saturated 
solution colors blue litmus tincture a wine-red color 
like carbonic acid. 

When boric acid is evaporated with an excess of 
concentrated phosphoric acid, and the dry residue 
treated with water in order to separate the phosphoric 



152 ELEMENTARY CHEMISTRY 

acid, a white amorphous mass is left, which possesses 
the composition BP0 4 . This substance is infusible, it 
is not attacked by strong acids, but dissolves in aque- 
ous potash. The existence of these compounds points 
to the conclusion that boron trioxide possesses feebly 
basic properties, resembling, in this respect, alumina, 
Al 2 O s , which also sometimes acts as a weak acid, 
and sometimes as a weak base, forming a correspond- 
ing phosphate, A1P0 4 . 

Hydrogen, Bromine and Oxygen Compounds. 
Hypobromus Acid, HBrO and Bromic Acid, HBr0 3 . 

Hypobromus Acid: Formula — HBrO. 

This acid and the corresponding salts, termed hypo- 
bromites, are formed, in a similar manner to hypo- 
chlorous acid, by the action of bromine on certain 
metallic oxides. Thus if bromine water be shaken up 
with mercuric oxide, and if the yellow liquid thus 
formed be treated successively with bromine and the 
oxide, a solution is obtained which contains in every 
100 c.c. 6.2 grams of bromine combined as hypobro- 
mous acid, the reaction being as follows, 

HgO+2Br 2 +H 2 0=2HBrO+HgBr 2 . 

The greater part of the hypobromous acid contained 
in this strong solution is decomposed on distillation 
into bromine and oxygen. It can, however, be dis- 
tilled in vacuo at a temperature of 40° without under- 
going this change. 

Aqueous hypobromous acid is a light straw-yellow 
colored liquid, closety resembling in its properties hy- 



HYDROGEN 153 

pochlorous acid, acting as a powerful oxidizing agent 
and bleaching organic coloring matters. 

If bromine be dropped very slowly into a cooled 
solution of an alkali hydroxide, a hypobromite is 
formed along with bromide, but it is very unstable, 
and changes quickly into bromate. 

By the action of bromine on lime, a substance simi- 
lar to bleaching powder is formed, and this salt was 
formerly termed bromide of lime. 

Bromic Acid: Formula — HBr0 3 . 

When bromine is dissolved in hot caustic potash or 
soda, a colorless solution is produced which contains a 
mixture of a bromide and a bromate, thus 

3Br 2 +6KHO=5KBr+KBr0 3 +3H 2 0. 
The sparingly soluble potassium bromate may be 
easily separated from the very soluble bromide by 
crystallization. Potassium bromate is also formed 
when bromine vapor is passed into a solution of po- 
tassium carbonate which has been ensaturated with 
chlorine gas. 

Solutions of alkali bromides are converted into bro- 
mates by electrolysis, the yield being almost quantita- 
tive if a little potassium chromate be added, in the 
case of the bromides of the alkaline earths some hypo- 
bromite is formed as well as bromate if the solution 
be kept cold. 

Free bromic acid is formed when chlorine is passed 
into bromine water, thus 

Br 2 +5Cl 2 +6H 2 O=2HBrO 3 +10HCl. 

The acid is, however, best obtained by the decom- 
position of the slightly soluble silver bromate. This 



154 ELEMENTARY CHEMISTRY 

salt is thrown down on the addition of silver nitrate 
to a solution of a soluble bromate, the precipitate thus 
prepared is well washed with water and then treated 
with bromine, bromic acid remains in solution and the 
insoluble silver bromide is thrown down, 

5AgBr0 3 +3Br 2 +3H 2 0=5AgBr+6HBr0 3 . 

Hydrochloric and hydriodic acids decompose bromic 
acid in a similar manner with formation of the chlor- 
ide or iodide of bromine. 

The bromates are as a rule sparingly soluble in 
water, and decompose on heating into oxygen and a 
bromide, but no per-bromate is formed in the process. 

Hydrogen, Nitrogen and Oxygen Compounds. 
Nitric Acid, HN0 3 and Nitrous Acid, HN0 2 . 

Nitric Acid or Hydrogen Nitrate: Formula — HN0 3 . 

Nitre, or potassium nitrate, is generally formed by 
the gradual oxidation of nitrogenous animal matter 
in presence of the alkali potash. Spring water, es- 
pecially the surface well-water of towns, frequently 
contains nitrates in solution, owing to water passing 
through soil containing decomposing animal matters, 
which by oxidation yield nitrates. For this reason, 
water containing nitrates is unfit for drinking pur- 
poses. Potassium nitrate, KN0 3 , commonly called salt- 
petre, occurs as an incrustation on the soil in various 
localities, especially in India, and sodium nitrate, 
NaNO.., or Chili saltpetre, is found in large beds on 
the coast of Chili and Peru. Nitric acid is obtained 
by heating nitre, KN0 3 , with sulphuric acid, or hydro- 
gen sulphate, H 2 S0 4 , when nitric acid, HN0 3 , and 



HYDROGEN 155 

hydrogen potassium sulphate, HKS0 4 , are formed. 
The decompositions here effected may serve as a type 
of a very large number of chemical changes classed 
as double decompositions. These may all be repre- 
sented as consisting in an exchange between two ele- 
ments, or groups of elements, thus, in the case in ques- 
tion, one atom of the hydrogen in sulphuric acid 
changes place with one atom or its equivalent of potas- 
sium in the nitre. These double decompositions may 
be represented in the form of an equation, in which 
one side signifies the arrangement and relative weights 
of the elements before combination, the other the ar- 
rangement and relative weights of the same elements 
after the chemical change has taken place, thus 

KN0 3 +H 2 S0 4 =HNO a +HKS0 4 

or, Nitre and Sulphuric Acid give Nitric Acid and Hy- 
drogen Potassium Sulphate. The relative weights of 
the elements and compounds entering into the de- 
composition are easily ascertained when we remember 
that each symbol expresses not merely the nature of the 
element, but also the relative weight w r ith which it 
combines, and that the combining weight of a com- 
pound is the sum of the combining weights of its con- 
stituents. The numbers expressed by the above equa- 
tion are 

K N O3+H2 S 0, = H N Os+H K S 0, 

39.1 + 14 + 48 + 2 + 32 + 64=1 + 14 + 48 + 1 + 39.1 + 32 + 64 
101.1 + 98 = 63 + 136.1 

We may express these double decompositions perhaps 



156 



ELEMENTARY CHEMISTRY 



more clearly if we represent the actual exchange of 
hydrogen for potassium, by a straight line, thus 

H | H S0 4 

N0 3 | K 

This signifies that, if we require 63 parts by weight 
of nitric acid, we shall require to take exactly 101.1 
parts of nitre and 98 parts of sulphuric acid, and that 
we shall have 136.1 parts of hydrogen potassium sul- 




Figf. 23. 

phate formed. Knowing these numbers, it is easy to 
calculate the proportions of ingredients needed to pro- 
duce any given quantity of nitric acid. 

Nitric acid is prepared on a small scale by placing 
about equal weights of nitre and sulphuric acid in a 
stoppered retort, which is gradually heated by a Bun- 
sen burner, as in Fig. 23. The nitric acid formed dis- 
tills over, and may be collected in a flask cooled with 
water. On a large scale this substance is prepared in 



HYDROGEN 157 

iron cylinders, into which the charges of nitre and acid 
are brought, the nitric acid being collected in large 
stoneware bottles. 

Nitric acid thus obtained is represented by the form- 
ula HN0 2 , it is a strongly fuming liquid, colorless when 
pure, but usually slightly yellow from the presence of 
lower oxides of nitrogen. In specific gravity it is 1.51 
at 18°, it does not possess a constant boiling point, as 
it gradually undergoes decomposition by boiling and 
becomes weaker. If mixed with water, and distilled 
under the ordinary atmospheric pressure, the residual 
acid is found at last to attain a fixed composition, boil- 
ing constantly at 120.5°, containing 68 per cent, of 
HN0 2 , and possessing a specific gravity of 1.414. When 
mixed with less water, a stronger acid than this comes 
over ; when mixed with more water, a weaker one first 
distills over till this constant composition is attained. 
Nitric acid contains 76 per cent, of oxygen, with some 
of which it easily parts, hence it acts as a strong oxidiz- 
ing agent. This is seen when we bring a small quantity 
(of metallic copper or tin into this liquid diluted with 
a little water, red fumes are immediately given off, and 
the metals are oxidized, for the same reason nitric acid 
bleaches indigo solution, oxidizing, and therefore de- 
stroying, the coloring matter. This reaction, and the 
formation of red fumes in presence of metallic copper, 
etc., serve as modes of detecting the presence of nitric 
acid. One of the most delicate tests for this acid con- 
sists in adding to the liquid to be tested an equal vol- 
ume of strong sulphuric acid, well cooling the mixture, 
and then carefully pouring on to its surface a solution 
of ferrous sulphate, FeS0 4 . A black ring is produced 
where the two layers of liquid meet if any nitric acid 



158 ELEMENTARY CHEMISTRY 

be present. Nitric acid forms, with metallic oxides, by 
the process of double decomposition, a numerous family 
of salts called nitrates. These are nearly all soluble 
in water, and many of them are largely used in the arts 
for various purposes. 

In nitric acid we have the first example of a series of 
important compounds known as acids. Most of the 
acids are soluble in water, they possess an acid taste, 
and have the property of turning blue litmus-solution 
red. All acids contain hydrogen, combined either with 
an element, or with a group of elements which almost 
always contains oxygen. These acids may be regarded 
as water in which parts of the hydrogen is replaced by 
the oxygenated group of atoms. When the rest of the 
hydrogen of an acid is replaced by a metal, as for in- 
stance when sulphuric acid acts upon zinc, the acid 
character of the substance disappears, and a salt, called 
zinc sulphate, is formed. Salts are likewise produced 
when certain hydroxides and oxides are brought into 
contact with acids, thus if the solution of potassium 
hydroxide (caustic potash), obtained by the action of 
the metal potassium on water, is added to nitric acid, 
the alkaline or caustic properties of the hydroxide as 
well as the sour taste of the nitric acid disappear at a 
certain point, the solution becomes neutral, that is, it 
does not change the color of either blue or red litmus, 
and the salt potassium nitrate is contained in the liquid. 
The soluble hydroxides which thus act upon acids are 
termed alkalies, and have the power of turning red lit- 
mus-solution blue. In the same way many metallic 
oxides, called basic oxides or bases, act upon acids to 
form salts, thus silver oxide dissolves in nitric acid, 



HYDROGEN 159 

and neutralizes its acid character, forming soluble sil- 
ver nitrate. 

Nitrous Acid: Formula — HN0 2 . 

Nitrogen trioxide dissolves in ice-cold water giving 
rise to a beautiful blue liquid, which contains nitrous 
acid. Nitrous acid is riot known in the pure state, it 
being a very unstable substance, which even in aqueous 
solution rapidly undergoes decomposition when 
warmed, giving rise to nitric acid and nitric oxide gas ; 
thus 

3HN0 2 =HN0 3 +2NO+H 2 0. 

This reaction is, however, a reversible one, nitric oxide 
yielding with nitric acid and water a small quantity 
of nitrous -acid, the condition of equilibrium being 
reached when about nine molecules of nitric acid are 
present to one of nitrous acid. The salts of this acid, 
or the nitrites, are, on the contrary, very stable bodies. 
They are not only formed by the action of the acid 
upon oxides, but also by the reduction of nitrates and 
by the oxidation of ammonia. Thus, for instance, po- 
tassium nitrate, KN0 2 , is formed either by fusing salt- 
petre, or, more easily, by heating this salt with lead 
or copper, thus 

2KN0 3 =2KX0 2 +0 2 . 

It is also formed to some extent by the action of sun- 
light on a sterilized solution of potassium nitrate. Ni- 
trites also occur in nature. Thus the atmosphere con- 
tains small quantities of ammonium nitrite, and traces 
of nitrites have been detected in the juices of certain 
plants. All the normal nitrites are soluble in water, 
and most of them soluble in alcohol. The silver salt is 



160 ELEMENTARY CHEMISTRY 

the nitrite which is most sparingly soluble in cold wa- 
ter, crystallizing out in long glittering needle-shaped 
crystals when the hot aqueous solution is cooled. 

When sodium nitrite is reduced by sodium amalgam, 
nitrogen and ammonia are the chief products, when the 
solution is hot, in the cold, nitrous oxide, hydroxyla- 
mine, and hyponitrous acid are formed, the relative 
amounts varying with the concentration. 

The nitrites deflagrate when thrown on to glowing 
carbon, as do the nitrates. They can, however, be dis- 
tinguished from the latter salts by the action of dilute 
acids, which produce an evolution of red fumes from 
the nitrites but not from the nitrates. In a similar way 
aqueous solutions of the neutral nitrites become of a 
light brown color when mixed with a solution of fer- 
rous sulphate, and this color deepens to a dark brown 
on the addition of acetic acid. In order to detect the 
presence of a nitrite in dilute solution in the absence 
of other oxidizing agents, iodide of potassium, starch 
paste, and dilute sulphuric acid are added. The latter 
acid sets free the nitrous acid, and this instantly de- 
composes the iodide with liberation of iodine. 

Nitrous acid is quantitatively oxidized by potassium 
per manganate and by hydrogen peroxide to form ni- 
tric acid, and is reduced to ammonia by boiling with 
ferrous sulphate and caustic soda. It may be esti- 
mated by any of these methods and also by its action 
on urea, which it oxidizes to nitrogen, carbon dioxide, 
and water. The free acid can, moreover, be titrated 
with caustic soda in dilute solution with methyl orange 
as indicator. 



HYDROGEN 161 

Hydrogen, Sulphur and Oxygen Compounds. 
Sulphuric Acid, H 2 S0 4 and Sulphurous Acid, H 2 S0 3 . 

Sulphuric Acid or Oil of Vitriol: Formula — H 2 S0 4 . 

This substance is the most important and useful acid 
known, as by its means nearly all the other acids are 
prepared, and also because it is very largely used in the 
arts and manufactures for a great variety of purposes. 
It has been truly said that the commercial prosperity 
of a country may be judged by the amount of sulphuric 
acid which it consumes. 

Sulphuric acid was first prepared by distilling a 
compound of iron, oxygen, sulphur, and water, called 
ferrous sulphate or green vitriol. The acid thus ob- 
tained is known as fuming acid, and consists of a mix- 
ture of hydrogen sulphate and sulphur trioxide, H 2 SO, t 
+SQ 2 . This plan of preparation has, however, long 
been superseded by the following more convenient 
method, which depends upon the fact that, although 
sulphur dioxide does not combine with free oxygen 
and water to form sulphuric acid, it is capable of tak- 
ing up the oxygen when the latter is united with nitro- 
gen in the form of nitrogen trioxide, N 2 3 . thus 
S0 2 +H 2 0+N 2 3 =H 2 SO,+N 2 2 " 

Sulphur dioxide, water and nitrogen trioxide yield 
sulphuric acid and nitric oxide. 

The nitric oxide formed in this decomposition takes 
up another atom of oxygen from the air, becoming 
N o o , and this is a^ain able to convert a second mole- 
cule of S0 2 with H 2 into H 2 S0 4 , being a second time 
reduced to N 2 2 , and ready again to take up another 



162 



ELEMENTARY CHEMISTRY 



atom of oxygen from the air. Hence it is clear that the 
acts simply as a carrier of oxygen between the 



N 2 2 

air and the S0 2 , an indefinitely small quantity of this 
nitrogen trioxide being, therefore, theoretically able to 
convert an indefinitely large quantity of sulphur diox- 
ide, water, and oxygen into sulphuric acid. 

This process is carried on, on the large scale, in 
chambers made of sheets of lead (often a capacity of 




Fig-. 24. 



5Q,000 or 100,000 cubic feet) supported upon wooden 
beams and uprights, into which the above-mentioned 
materials are brought. Fig. 24 shows the arrangements 
for the manufacture of sulphuric acid. The leaden 
chambers, of which two are represented in the figure, 
are connected together by a wide leaden passage, and 
the gases passing from the first into the second cham- 
ber thus become thoroughly mixed. The sulphur diox- 



HYDROGEN 183 

ide is procured either by burning sulphur in a current 
of air, or by roasting a mineral called iron pyrites, a 
compound of sulphur and iron, FeS 2 , in a suitable 
furnace aa. The sulphur of the pyrites burns away, 
and the gaseous product is led, together with atmos- 
pheric air, into the chamber, whilst ferric oxide, Fe 2 3 , 
remains behind in the furnace. A small stove b con- 
taining nitre, is placed in the central part of the fur- 
nace, where this salt is decomposed by the action of 
the sulphur dioxide, an alkaline sulphate being formed, 
whilst nitrous fumes pass with the other gases into the 
chamber. Jets of steam are also blown into the cham- 
ber at various points, from a boiler c and a thorough 
draft is maintained by connecting the end of the cham- 
ber with a high chimney not shown in the figure, but 
placed beyond the tower d. The fumes, gases, and air 
escaping from the chamber have to pass through the 
tower d and there meet with a jet of steam, by means of 
which the soluble acid vapors are nearly all condensed 
before reaching the chimney. The sulphuric acid, as it 
forms, falls on to the floor of the chamber, and, when 
the process is working properly, it is continually drawn 
off, attaining a specific gravity of about 1.60, the waste 
gases passing out of the chamber should contain noth- 
ing but nitrogen and small quantities of nitric oxide. 
In order to obtain from this weak chamber acid the 
pure sulphuric acid, H 2 S0 4 , the excess of water must 
be removed by evaporation: this is conducted, on the 
large scale, first, by heating the chamber acid in cov- 
ered leaden pans e, until the specific gravity rises to 
1.72, when the acid is known as the brown oil of vitriol 
of commerce, and then further concentrated in vessels 
of glass, or of platinum (as lead is attacked by the 



164 ELEMENTARY CHEMISTRY 

strong acid), until its maximum strength and specific 
gravity is attained. The hydrogen sulphate thus ob- 
tained is a thick oily liquid boiling about 338 °, and 
freezing at 10.8°, its specific gravity at 0° is 1.854. It 
combines with water with great force, absorbing mois- 
ture rapidly from the air. Hence it is used in the 
laboratory as a drying agent. Great heat is evolved 
when this acid is mixed with water, and care must be 
taken to bring these two liquids together gradually, 
otherwise an explosive combination may ensue. Many 
organic bodies, such as woody fibre and sugar, are com- 
pletely decomposed and charred by strong sulphuric 
acid, whilst others, such as alcohol, oxalic and formic 
acid, are split up into other compounds by the with- 
drawal of the elements of water by this acid. 

One molecule of hydrogen sulphate unites with one 
of water to form a compound, H 2 S0 4 -f-H 2 0, which can 
be obtained pure by cooling a mixture of acid and 
water having a specific gravity of 1.78 down to 7° 
Centigrade, at which temperature rhombic crystals of 
the hydrated acid are formed. The sulphuric acid of 
commerce frequently contains large quantities of im- 
purities, especially lead sulphate from the chamber, 
and frequently arsenic from the pyrites, and nitric acid, 
as well as the lower oxides of nitrogen. In order to 
free the acid from these impurities it must be distilled 
and subjected to other treatment, for a description of 
which the reader is referred to the larger treatises. 
At high temperatures sulphuric acid decomposes into 
sulphur dioxide, S0 2 , oxygen, 0, and water, H 2 0. If 
a current of the acid be allowed to flew onto red-hot 
bricks, and the gases resulting from the decomposition 
passed through water, the sulphur dioxide will be com- 



HYDROGEN 165 

pletely absorbed, and a supply of pure oxygen ob- 
tained. Hydrogen sulphate is a dibasic acid, it contains 
two atoms of hydrogen, either or both of which can 
be replaced by an equivalent quantity of a metal. As 
with sulphurous acid, in the case of the alkaline metals, 
we have two salts, KHS0 4 and K 2 S0 4 . Barium and 
lead sulphates are insoluble in water, hence soluble 
salts of these metals are used as tests of the presence 
of a sulphate, a few drops of solution of barium chlo- 
ride, for example, producing an immediate white pre- 
cipitate of barium sulphate in water containing the 
merest trace of sulphuric acid or of a soluble sulphate. 
Calcium, strontium, and potassium-sulphates are but 
slightly soluble in water, whilst the other sulphates are 
easily soluble. 

Some sulphates crystallize as anhydrous salts, such 
as K 2 S0 4 , potassium sulphate, BaS0 4 , barium sulphate, 
and Ag 2 S0 4 , silver sulphate, whilst others require wa- 
ter to retain their crystalline form, and this water is 
termed water of crystallization. The crystals of iron 
sulphate or green vitriol, and of zinc sulphate or white 
vitriol, contain seven molecules of water in the solid 
form, while copper sulphate or blue vitriol requires but 
five molecules to preserve its crystalline form. 

Sulphurous Acid: Formula — H 2 S0 3 . 

This substance, like many other acids whose corre- 
sponding anhydrides are gaseous, is only known in 
aqueous solution. This solution smells and tastes like 
the gas and has a strongly acid reaction. Exposed to 
the light it is decomposed with formation of pentathi- 
onic acid. When an aqueous solution saturated at 3° 
is allowed to stand, crystals of a hydrate, H 2 S0 3 + 



166 ELEMENTARY CHEMISTRY 

6H 2 0, are obtained, and other hydrates of the formula 
H 2 S0 3 , 8H 2 0, H 2 S0 3 , 10H 2 O, and H 2 S0 3 , 14H 2 have 
been described. Sulphurous acid differs from the acids 
which have hitherto been described, inasmuch as it 
contains two atoms of hydrogen, both of which may 
be replaced by metals. It is therefore termed a dibasic 
acid, it forms two series of salts termed sulphites, in 
one of which only half of the hydrogen is replaced by 
a metal, and which may therefore be considered as be- 
ing at once a salt and a monobasic acid, another in 
which the whole of the hydrogen of the acid has been 
replaced by a metal. The salts of the first series are 
termed acid sulphites and of the latter normal sul- 
phites. 

The acid sulphites of potassium and sodium are ob- 
tained by passing sulphur dioxide gas into caustic soda 
or caustic potash as long as it is absorbed. If, then, 
exactly the same quantity of alkali is added to this 
solution as was originally taken for the preparation, 
the normal salts are obtained. All the sulphites of the 
alkali metals are easily soluble in water, the normal 
sulphites of the other metals being either sparingly 
soluble or insoluble in water. They dissolve, however, 
in aqueous sulphurous acid and exist in such a solution 
as acid salts, but on evaporation they decompose with 
formation of the normal salt and sulphurous acid. The 
normal sulphites have no odor, and those which are 
soluble in water possess a sharp taste. They are read- 
ily detected by the fact that when they are mixed with 
dilute sulphuric acid they give off sulphur dioxide, and 
also that their neutral solutions give a precipitate with 
barium chloride which is soluble in dilute hydrochloric 
acid, whereas if nitric acid be added to this solution 



IRON 167 

and the mixture warmed, a precipitate of barium sul- 
phate, formed by oxidation from the sulphite, is thrown 
down. 

IRON. 

Symbol — Fe. Atomic Weight — 55.6. 

Of all metals the most important to mankind, its uses 
were long unknown to the human race. The age of iron 
implements being preceded by those of bronze and 
stone. Pure metallic iron exists only in very small 
quantity on the earth's surface, almost entirely occur- 
ring in those peculiar structures known as meteoric 
stones, which possess an extra-terrestrial origin. 

The process of obtaining iron from its ores is a some- 
what difficult one, and requires an amount of knowl- 
edge and skill which the early races of men did not 
possess. The iron of commerce exists in three different 
forms, exhibiting very different properties, and pos- 
sessing different chemical constitutions : Wrought 
iron, Cast iron and Steel. 

The first is nearly pure iron, the second is a com- 
pound of iron with varying quantities of carbon and 
silicon, and the third a compound of iron with less car- 
bon than that needed to form cast iron. The modes of 
manufacture of these three kinds of iron are essential- 
ly different, and will be best understood when the prop- 
erties of the metal have been described. 

Pure iron in the form of powder may be obtained 
by reducing the oxide, moderately heated in a current 
of hydrogen, it must, however, be retained in an at- 
mosphere of hydrogen, as finely-divided iron takes fire 
and burns to oxide when exposed to the air. A button 
of pure iron may be prepared by exposing fine iron 



168 ELEMENTARY CHEMISTRY 

wire mixed with some oxide of iron to a very high 
temperature in a covered crucible, the oxide retaining 
the traces of impurity which the wire contained. Iron 
has a bright white color, and is remarkably tough, 
though soft, an iron wire two mm. in thickness not 
breaking until weighted with 250 kilogs. The pure 
metal crystallizes in cubes. Iron which has been uni- 
formly hammered exhibits, when broken, a granular 
and crystalline structure, this structure becomes, how- 
ever, fibrous when the iron is rolled into bars, and the 
more or less perfect form of the fibre determines to a 
great extent the value of the metal. This fibrous tex- 
ture of hammered bar iron undergoes a change when 
exposed to long-continued vibration, the iron return- 
ing to its original crystalline condition, and many ac- 
cidents have occurred in the sudden snapping of axles, 
owing to this change from the fibrous to the granular 
texture. Wrought iron melts at a very high tempera- 
ture, but as it becomes soft at a much lower point, it 
can be easily worked, as, when hot, it possesses the 
peculiar property of welding, that is, the power of 
uniting firmly when two clean surfaces of hot metal are 
hammered together. 

Iron and certain of its compounds are strongly mag- 
netic, but the metal loses this power when red hot, 
regaining it upon cooling. A solid mass of iron does 
not oxidize or tarnish in dry air, at the ordinary tem- 
perature, although iron powder takes fire spontaneous- 
ly, but if heated it oxidizes, with the production of 
black scales of oxide, and when more strongly heated 
in the air, or plunged into oxygen gas, it burns, with 
the formation of the same black oxide. In pure water 
iron does not lose its brilliancy, but if a trace of car- 



IRON 169 

bonic acid is present, and access of air is permitted, 
the iron begins at once to oxidize at the surface, or to 
rust, forming a hydrated sesquioxide. 

Iron and Oxygen Compounds. 

Iron Monoxide, FeO, Iron Sesquioxide or Ferric Oxide, 

Fe 2 A 3 and the Magnetic or Black Oxide 

of Iron, Fe 3 4 . 

Iron Monoxide: Formula — FeO. 

This substance has not been prepared in the pure 
state, owing to the readiness with which it absorbs 
oxygen, passing into the higher oxides. Hydrated fer- 
rous oxide, FeH 2 2 , is thrown down as a white pre* 
cipitate, when potash or soda is added to a soluble 
ferrous salt, this white precipitate can only be obtained 
in complete absence of oxygen, as it at once absorbs 
this gas, yielding a greenish-brown precipitate of a 
higher oxide. This oxide colors glass green, and gives 
the peculiar tint to common bottle-glass. 

Iron Sesquioxide or Ferric Oxide: Formula — Fe 2 3 . 

This oxide occurs native, as the minerals red haema- 
tite and specular iron ore, while combined with water, 
it forms brown hsematite. It may be readily prepared 
artificially by heating ferrous sulphate to redness, or 
by adding solution of ammonia or caustic potash to a 
solution of a ferric salt, when the hydrated oxide falls 
down as a bulky brownish-red powder, which dissolves 
in acids, forming the ferric salts, when thus acted upon 
by sulphuric acid, ferric sulphate, Fe 2 3S0 4 , is pro- 
duced, and by hydrochloric acid, ferric chloride, 
Fe 2 C] 6 . 



170 ELEMENTARY CHEMISTRY 

Black Oxide of Iron : Formula — Fe 3 4 . 

The magnetic oxide of iron occurs native, crystal- 
lized in octahedra, and, as the mineral loadstone, it 
constitutes one of the most valued ores of iron. It is 
the oxide formed when iron is oxidized at a high tem- 
perature in the air, in oxygen, or in aqueous vapor. 
A corresponding sulphide, Fe 3 S 4 , is also magnetic. 

The Manufacture of Iron. 

The oldest method of manufacturing wrought iron 
was to reduce it at once from the ore by heating in a 
furnace with charcoal or coal, and to hammer out the 
spongy mass of iron thus obtained. This plan can only 
be economically employed on a small scale and with 
the purest forms of iron ore, and has been superseded 
by a more complicated method, applicable, however, 
to all kinds of iron ore. This consists in the formation 
of cast iron as the first product, and the subsequent 
separation of the carbon and silicon which the east iron 
contains. Cast iron is manufactured chiefly from clay 
ironstone, which generally occurs in masses, situated 
in the immediate neighborhood of a coal seam. The 
ironstone is first roasted, in which operation the car- 
bonic acid is driven off, and ferric oxide formed, the 
ore afterwards being thrown, together with coal and 
limestone, into a blast furnace. It has the shape of a 
double cone, and is about fifty feet in height, and fif- 
teen to eighteen feet in width at the broadest part. 
The furnace is closed at the bottom, the air necessary 
for the maintenance of the combustion being supplied 
in a powerful blast, blown through pipes called tuyeres. 
The mixture of fuel and ore, being put in at the top of 



IRON 171 

the furnace, is added continually as the burning mass 
sinks down and the molten mass is drawn off at the 
bottom, so that one furnace often does not stop work- 
ing for several years. At the lowest part of the struc- 
ture is the hearth, where the melted metal and fused 
slag collect, the former being occasionally tapped from 
the bottom of the hearth, and cast into pigs in moulds 
made in the sand, while the lighter slag, which swims 
on the surface of the metal, runs continually out from 
an opening at the upper part of the hearth. 

The first chemical change which the roasted iron 
ore, or impure ferric oxide, undergoes in its passage 
from the top to the bottom of the furnace, is its reduc- 
tion to a porous mass of metallic iron, by the carbonic 
oxide gas proceeding from the lower layers of burn- 
ing coal. The temperature of this portion of the fur- 
nace is, however, much too low to melt the iron, and 
it therefore sinks down unchanged, together with the 
limestone, until it reaches a point at which the heat is 
greater. Here the second change occurs, the sand, and 
other impurities of the ore unite with the limestone 
to form a fusible silicate or slag, whilst the heated met- 
al, coming in contact with carbon, unites at once with 
it to form cast iron, a fusible compound, which runs 
down to the bottom of the furnace. This, in passing 
through the hottest portion of the furnace, reduces the 
silica, with which it meets, to silicon, and, combined 
with this, it forms cast iron. 

The properties and appearance of cast irons vary 
much with the quantity of carbon and silicon which 
they contain, for cast iron is not a definite chemical 
compound of these elements with iron. Carbon is found 
in cast iron, as scales of graphite, giving rise to mot- 



172 ELEMENTARY CHEMISTRY 

tied cast iron, and in combination, forming white cast 
iron. Sometimes sulphur and phosphorus are also 
found in cast iron, but these must be considered as im- 
purities. A great saving of fuel in the working of 
blast furnaces has been effected by employing the heat 
of combustion of the waste gases, which usually escape 
and burn at the top of the furnace, to raise the tem- 
perature of the blast of air supplying the furnace. 
The gases are collected at the top of the furnace by 
a hood and pass down an iron pipe, which is carried 
down to the furnaces in which the gases are burnt. 

In order to obtain wrought from cast iron, the latter 
must undergo the processes of refining and puddling. 
These consist essentially in burning out the carbon, 
silicon, sulphur, and phosphorus, by exposing the heat- 
ed metal to a current of air in a reverberatory furnace. 
The melted cast iron becomes first covered with a coat 
of oxide, and gradually thickens so as to allow of its 
being rolled into large lumps or balls. During this 
process the whole of the carbon escapes as carbonic 
oxide, and the silicon becomes oxidized to silica, which 
unites with the oxide of iron, and forms a fusible slag, 
any phosphorus or sulphur contained in the pig iron 
is also oxidized in this process. The ball is then ham- 
mered to give the metal coherence, and to squeeze out 
the liquid slag, and the mass is afterwards rolled into 
bars or plates. 

Another interesting branch of the iron, trade is the 
manufacture of steel. This substance is formed when 
bars of wrought iron are heated to redness for some 
time in contact with charcoal, the bar is then found to 
have become fine-grained instead of fibrous, the sub- 
stance is more malleable and more easily fusible than 



IRON 173 

the original bar iron, and is found to contain carbon 
varying -in amount from one to two per cent. Steel 
possesses several important properties, especially the 
power of becoming very hard and brittle when quickly 
cooled, which fits it for the preparation of cutting- 
tools, these are, however, generally made of bar-steel, 
which has been previously fused and cast into ingots. 

A new and very rapid mode of preparing steel, which 
is of high industrial importance, is that known as the 
Bessemer process. This process consists in burning out 
all the carbon and silicon in cast iron by passing a 
blast of atmospheric air through the molten metal, and 
then in adding such a quantity of a pure cast iron to 
the wrought iron thus prepared as is necessary to give 
carbon enough to convert the whole mass into steel, 
the melted steel is then at once cast into ingots. In this 
way six tons of cast iron can at one operation be con- 
verted into steel in twenty minutes. 

Iron, Carbon and Oxygen Compounds. 
Carbonate of Iron : Formula — FeC0 3 . 

Ferrous Carbonate is an insoluble compound, and 
occurs largely as a mineral called spathose iron ore, 
which is isomorphous with calc-spar. It also occurs in 
a less pure form, constituting the clay ironstone, the 
ore of iron from which a large proportion of iron is 
prepared. 

Iron and Chlorine Compounds. 
Ferrous Chloride: Formula — FeCl 2 . 

When dry hydrochloric acid gas is passed over hot 
metallic iron, ferrous chloride and hydrogen are 
formed. The hydrated chloride is also produced when 



174 ELEMENTARY CHEMISTRY 

iron is dissolved in aqueous hydrochloric acid, green 
crystals being deposited. 

Iron, Sulphur and Oxygen Compounds. 
Sulphate of Iron: Formula— FeS0 4 . 

Ferrous Sulphate, sometimes called green vitriol, is 
obtained by dissolving metallic iron, or ferrous sul- 
phide, in sulphuric acid; and is also prepared by the 
slow oxidation of pyrites. 

The solution thus obtained yields on evaporation 
large green crystals of the salt. It is largely used in 
the manufacture of several black dyes, and is one of 
the constituents of writing-ink. Like all the ferrous 
compounds, this salt easily takes up oxygen, produc- 
ing a new salt called ferric sulphate. 

Iron and Sulphur Compounds. 
Ferrous Sulphide: Formula — FeS. 

Ferrous Sulphide is an invaluable compound, formed 
by fusing equivalent quantities of sulphur and iron to- 
gether, it is employed in the laboratory for the genera- 
tion of sulphuretted hydrogen. A disulphide, called 
iron pyrites, is found in large quantities, and is much 
used in the production of sulphuric acid. 

IODINE. 

Symbol— I. Atomic Weight — 125.9. 

Iodine occurs combined with metals in sea-water, and 
is obtained from kelp, the ash of certain sea-weeds, in 
which it is found as the iodides of sodium and mag- 
nesium. Iodine is obtained from kelp by exactly the 
same process as that by which clilorine and bromine 



IODINE 175 

are obtained from chlorides and bromides, by heating 
with sulphuric acid and manganese dioxide. Iodine 
is thus liberated in the form of a deep violet-colored 
vapor, which condenses to a dark grey solid, with 
bright metallic lustre. Iodine melts at 115°, and boils 
above 200°, and has a specific gravity of 4.95. It gives 
off a perceptible amount of vapor at the ordinary tem- 
perature, and possesses a faint chlorine-like smell. 
Water dissolves a very small quantity of iodine, but 
in presence of a soluble iodide it is freely dissolved, 
forming a deep red or brown solution. It is easily 
soluble in alcohol, giving a reddish-brown solution, and 
in carbon disulphide and chloroform, imparting to 
them a splendid violet color. Iodine does not possess 
such active properties as either of the preceding ele- 
ments, its solution does not bleach organic coloring 
matters, and it is liberated from its compounds by both 
bromine and chlorine. Free iodine forms a remarkable 
compound with starch, of a splendid blue color, and by 
this means the minutest trace of this substance can be 
detected. To apply this test, one drop of potassium 
iodide solution is added to starch paste largely diluted 
with water, no blue color is observed until the iodine 
is set free by the addition of a drop or two of chlorine- 
water, when a deep blue coloration is instantly per- 
ceived. Iodine acts as a powerful poison, but, given 
in small quantities, it is much used as medicine. 

The crude Chili saltpetre forms the chief source of 
iodine, to obtain the iodine the last mother-liquor ob- 
tained in the preparation of the sodium nitrate, which 
contains about 20 per cent of sodium iodate, is treated 
with a solution of sodium bisulphite, the iodine sepa- 
rates out in the solid form, and is filtered off and puri- 



176 



ELEMENTARY CHEMISTRY 



lied by resublimation, the vapors being condensed in 
a series of udells similar to those shown in Pig. 25. 

In order to purify the commercial iodine it is washed 
with a small quantity of water, dried on porous plates, 
and resublimed. The only mode of obtaining chem- 
ically-pure iodine, free from every trace of chlorine 
and bromine, is to dissolve the commercial resublimed 
substance in iodide of potassium solution, and then to 




Fig. 25. 



precipitate the iodine by water. The precipitate is well 
washed with water and then distilled with steam, the 
solid iodine in the distillate collected and dried in 
vacuo over solid nitrate of calcium, which is frequent- 
ly changed, and distilled afterwards over solid caustic 
baryta to remove the last traces of water and of hy- 
driodic acid. 

Iodine is a bright, shining, crystalline, blackish-grey 



IODINE 177 

solid, which is usually opaque, but may be obtained 
in transparent films by deposition on glass surfaces at 
a temperature of about — 180°. The crystals when 
large possess almost a metallic lustre, it crystallizes by 
sublimation in the form of prisms or pyramids. 

In its chemical properties iodine resembles chlorine 
and bromine, the two latter elements have the power 
of displacing iodine from its combination with metals, 
or electro-positive elements, thus : 

2KI+C1 2 =2KC1+I 2 . 

The compounds of iodine with oxygen, or with electro- 
negative elements are, on the other hand, more stable 
than those of the other two elements. Thus iodine ex- 
pels chlorine from the chlorates with formation of 
iodate and free chlorine, 

2KC10 3 +I 2 =2KI0 3 +C1 2 . 

Iodine is very sparingly soluble in water, 1 part dis- 
solving in 3,750 parts of w-ater at 15° and in 2,200 at 
30°. It dissolves readily in a solution of potassium 
iodide, forming a brown solution, in which, however, 
the iodine is probably not present as such, but as the 
unstable potassium triiodide. It also dissolves readily 
in a large number of other solvents, in some of which, 
such as chloroform, carbon bisulphide, and liquid hy- 
dro-carbons, it gives rise to violet solutions, whilst with 
others, such as alcohol, ether, organic acids and esters, 
and pyridine, it yields yellow or brown solutions. The 
cause of the difference of color of the various solutions 
has not yet been ascertained w r ith certainty, but it 
seems most probable that in the violet solutions iodine 
is present as such, whilst the brown solutions contain 
an unstable compound of iodine with the solvent. The 



178 ELEMENTARY CHEMISTRY 

first-named class of solutions when concentrated are 
black and opaque, in the case of carbon bisulphide, 
such a solution is diathermanous, allowing the invisible 
heating rays of low refrangibility to pass, but not the 
visible rays. 

Although iodine, unlike chlorine and bromine, does 
not combine readily with hydrogen, it unites with many 
of the metals and non-metals with evolution of light 
and heat. Thus solid phosphorus, when brought into 
contact with iodine, first melts and then bursts into 
flame owing to the heat evolved in the act of combina- 
tion, and powdered antimony takes fire when thrown 
into iodine vapor, antimony iodide being produced, 
while if the vapor of mercury be passed over heated 
iodine, immediate action occurs, the iodides of mercury 
being formed. When iodine is brought into contact 
with water and filings of iron or zinc, a violent reac- 
tion occurs, colorless solutions of the respective iodides 
resulting. The action of iodine upon the alkali-metals 
is analogous to that of chlorine and bromine. Sodium 
and iodine can be heated together without alteration, 
whilst if potassium be employed an explosive combina- 
tion occurs. 

LEAD. 

Symbol— Pb. Atomic Weight— 204.85. 

Lead does not usually occur free in nature, all the 
lead of commerce is obtained from galena, or lead sul- 
phide. The mode of reducing lead from this ore is a 
very simple one, the galena is roasted in a reverbera- 
tory furnace, with the addition of a small quantity of 
lime to form a fusible slag with any silicious mineral 
matter present in the ore. By the action of the air a 



LEAD 179 

portion of the sulphide is oxidized to sulphate, whilst 
in another portion the sulphur burns off as sulphur 
dioxide, and lead oxide is left behind, after the lapse 
of a certain time the air is excluded and the heat of 
the furnace raised, the lead sulphate and oxide formed 
both decompose the remaining sulphide, giving off sul- 
phur dioxide and leaving metallic lead behind. 

Galena almost always contains small quantities of 
silver. Lead is a bluish-white colored metal, and so 
soft that it may be scratched with the nail, it may be 
drawn out to wire, or hammered into plate, but pos- 
sesses little tenacity or elasticity, and a wire 2 mms. in 
diameter breaks with a load of 2 kilos. Lead melts at 
334°, and at a higher temperature volatilizes, though 
not in quantity sufficient to enable it to be distilled. 

The bright surface of the metal remains permanent- 
ly in dry air, but it soon becomes tarnished in moist 
air, owing to the formation of a film of oxide, and this 
oxidation proceeds rapidly in presence of a small quan- 
tity of weak acid, such as carbonic or acetic. In pure 
water freed from air lead also preserves its lustre, but 
if air be present, lead-oxide is formed, and this dis- 
solving slightly in the water a fresh portion of metal 
is exposed for oxidation. This solvent action of water 
upon lead is a matter of much importance, owing to 
the common use of lead water-pipes, and the peculiarly 
poisonous action of lead compounds upon the system 
when taken even in minute quantities for a length of 
time. The small quantity of certain salts contained 
in all spring and river waters exerts an important in- 
fluence on the action of lead. The waters containing 
nitrates or chlorides are liable to contamination with 
lead, while those hard waters containing sulphates or 



180 ELEMENTARY CHEMISTRY 

carbonates may generally be brought into contact with 
lead without danger, as a thin deposit of sulphate or 
carbonate is formed, which preserves the metal from 
further action. If the water contains much free car- 
bonic acid, it should not be allowed to come into con- 
tact with lead, as the carbonate dissolves in water con- 
taining this substance. The presence of lead in watei* 
may easily be demonstrated by passing a current of 
sulphuretted hydrogen through a deep column of the 
acidified water, and noticing whether the liquid be- 
comes tinged of a brown color, owing to the formation 
of lead sulphide. 

Lead and Chlorine Compounds. 
Lead Chloride: Formula — PbCl 2 , 

Lead Chloride is prepared by adding hydrochloric 
acid to a strong solution of lead nitrate, when a crys- 
talline precipitate of lead chloride is formed. It dis- 
solves in about thirty parts of boiling water, separating 
out in shining needles on cooling. 

Lead and Chromium and Oxygen Compounds. 
Lead Chromate: Formula — PbCr0 4 . 

Lead Chromate is a yellow insoluble salt, used as a 
pigment under the name of chrome-yellow. 

Lead and Iodine Compounds. 
Lead Iodide: Formula — Pbl 2 . 

Lead Iodide is precipitated in the form of splendid 
yellow spangles, when hot solutions of potassium iodide 
and lead nitrate are mixed and allowed to cool. 



LEAD 181 

Lead and Oxygen Compounds. 

Lead Monoxide, PbO, Lead Dioxide, Pb0 2 and Red 
Oxide or Red Lead. 

Lead Monoxide or Litharge: Formula — PbO. 

Lead Monoxide, or Litharge, is a straw-colored pow- 
der, obtained by heating lead in a current of air, it 
fuses at a red heat, forming scaly crystals termed 
Litharge or Massicot. Lead monoxide is soluble in 
caustic potash, and is deposited from a hot solution in 
the form of rhombic prisms. This oxide forms with 
acids the important series of lead salts, which are gen- 
erally colorless, and of which the soluble ones act as 
violent poisons. Lead monoxide combines with silica 
to form an easily fusible silicate, or glass. Earthen 
crucibles in which the oxide is fused are rapidly at- 
tacked. A white hydrated oxide is obtained by pre- 
cipitating a soluble salt of lead by caustic potash, and 
this if heated yields the oxide. 

Lead Dioxide : Formula — Pb0 2 . 

Lead Dioxide is a brown powder obtained by pass- 
ing chlorine through the hydrated monoxide, or by 
digesting red lead with nitric acid. Lead dioxide does 
not form salts with acids. When heated it yields half 
its oxygen, and acted upon with warm hydrochloric 
acid, chlorine is evolved, and lead chloride is formed. 

Lead Oxide: 

Red Oxide, or Red Lead, a compound of the two last 
oxides, having the composition 2PbO+Pb0 2 . It is ob- 
tained by exposing massicot to the air at a moderate 



182 ELEMENTARY CHEMISTRY 

red heat, oxygen being absorbed. Red lead is chiefly 
used in glass-making and steam-fitting. When treated 
with dilute nitric acid the lead monoxide dissolves, 
forming soluble lead nitrate, leaving the puce-colored 
oxide behind. 

Lead, Sulphur and Oxygen Compounds. 
Lead Sulphate : Formula — PbS0 4 . 

Lead Sulphate is a white insoluble salt, which is 
found native, and is prepared artificially by adding 
sulphuric acid to a soluble lead salt. 

LITHIUM. 
Symbol— Li. Atomic Weight— 6.98. 

This metal is prepared by decomposing the fused 
chloride by electricity, it is of a white color, it fuses 
at 180°, and is the lightest metal known. The lithium 
salts were formerly supposed to be very rare, only be- 
ing known to occur in three or four minerals, but spec- 
trum analysis has shown that this is a widely-distrib- 
uted substance. It occurs in small quantities in almost 
all waters, in milk, tobacco, and even in human blood. 
Lithium in its chemical relations stands between the 
class of alkaline and alkaline-earth metals, the hydrate, 
carbonate, and phosphate being only sparingly soluble 
in water. All the volatile lithium compounds impart a 
magnificent crimson tinge to the flame, and the spec- 
trum of this flame exhibits the presence of one bright 
and very characteristic red line, by means of which 
the presence of the minutest trace of this substance 
can be detected with certainty and ease. 



MAGNESIUM 183 

MAGNESIUM. 

Symbol— Mg. Atomic Weight— 24.2. 

This metal occurs in large quantities as carbonate, 
along with calcium carbonate, in mountain limestone, 
and also in sea-water and certain mineral springs, as 
chloride and sulphate. The metal itself has only re- 
cently been prepared in quantity, it is best obtained 
by heating magnesium chloride with metallic sodium, 
sodium chloride and metallic magnesium being formed. 
This metal is of a silver-white color, and fuses at a low 
red heat, it is volatile, and may be easily distilled at a 
bright red heat, when soft it can be pressed into wire, 
and with care it may be cast like brass, although when 
strongly heated in the air it takes fire and burns with 
a dazzling white light, with the formation of its only 
oxide, magnesia. The light emitted by burning mag- 
nesium wire is distinguished for its richness in chem- 
ically active rays, and this substance is therefore em- 
ployed as a substitute for sunlight in photography, and 
has been employed with success for photographing the 
interior of caverns. 

Magnesium does not oxidize in dry air, it is only 
slowly acted upon by cold water, but more rapidly by 
hot water, it rapidly dissolves in sulphuric and hydro- 
chloric acids, with evolution of hydrogen. 

Magnesium, Carbon and Oxygen Compounds. 
Magnesium Carbonate.: Formula — MgC0 3 . 

Magnesium Carbonate is an insoluble compound, oc- 
curring as a crystallized mineral termed magnesite. 
The white magnesia is a varying mixture of carbonate 



184 ELEMENTARY CHEMISTRY 

and hydrate, made by precipitating a hot solution of 
magnesium sulphate with sodium carbonate. Mag- 
nesium sulphide is not formed in the wet way. Mag- 
nesium resembles in many respects the metals of the 
alkaline earths, but it may be distinguished from these 
by the solubility of the carbonate in ammonium chlo- 
ride, as well as by the ready solubility of the sulphate 
in water. Magnesium forms an insoluble double phos- 
phate with ammonia, and it is in this form the metal 
is usually estimated. 

Magnesium and Chlorine Compounds. 
Magnesium Chloride: Formula — MgCl 2 . 

Magnesium Chloride is a fusible salt obtained by 
evaporating magnesia dissolved in hydrochloric acid 
with an equal quantity of sal-ammoniac, on fusion, the 
latter salt volatilizes, and the magnesium chloride re- 
mains behind. 

Magnesium and Oxygen Compounds. 
Magnesium Oxide or Magnesia: Formula — MgO. 

Magnesia is a light white amorphous infusible pow- 
der, obtained by heating the carbonate or nitrate, and 
is largely used in medicine, and known as calcined 
magnesia. It unites with acids to form the magnesium 
salts, but it does not possess strong alkaline reaction. 

Magnesium, Sulphur and Oxygen Compounds. 
Magnesium Sulphate : Formula — MgS0 4 +7H 2 0. 

Magnesium Sulphate is a soluble substance known as 
Epsom Salts, it occurs in spring water, and contains 
seven atoms of water of crystallization, it is now large- 



MANGANESE 185 

ly made from dolomite by separating the lime with the 
sulphuric aeid. Magnesium sulphate forms, with the 
alkaline sulphates, double salts, in which the alkaline 
sulphate takes the place of one molecule of the water 
of crystallization. 

MANGANESE. 
Symbol— Mn. Atomic Weight— 54.6. 

Manganese occurs in nature as an oxide, and it can 
be obtained, though with difficulty, in the metallic state 
by heating the oxide very strongly with charcoal. The 
metal is of a reddish-white color, it is brittle, and hard 
enough to scratch glass. It decomposes water at the 
ordinary temperature, with evolution of hydrogen, it 
cannot be preserved in the air without undergoing oxi- 
dation, and must be kept under naphtha, or in a sealed 
tube, it is slightly magnetic, and, like iron, combines 
with carbon and silicon. Metallic manganese is not 
used in the arts, but an alloy of this metal and iron is 
now made on a large scale, and used in the manufac- 
ture of steel. Some of its oxides are used for the pur- 
pose of evolving chlorine from hydrochloric acid, and 
also for tinting glass a purple color. 

Manganese and Oxygen Compounds. 

Manganese Monoxide, MnO, Manganese Sesquioxide, 
Mn 2 3 and Manganese Dioxide, Mn0 2 . 

Manganese Monoxide: Formula — MnO. 

Manganese Monoxide is a greenish powder, obtained 
by heating the carbonate in absence of air, it forms 
with acids a series of pink-colored salts, and rapidly 



186 ELEMENTARY CHEMISTRY 

absorbs oxygen, passing into a higher state of oxida- 
tion. The hydrate is precipitated as a white gelatinous 
mass, when an alkali is added to a solution of man- 
ganous salt, this, however, rapidly becomes brown, 
owing .to absorption of oxygen. 

Manganese Sesquioxide : Formula — Mn 2 3 . 

Manganese Sesquioxide exists in nature as braunite, 
and may be prepared artificially by exposing man- 
ganous oxide to a red heat. It forms a series of some- 
what unstable salts, of which the manganese alum is 
one of the most interesting, being isomorphous with 
common alum, in which Mn 2 3 is substituted for A1 2 3 . 

Manganese Dioxide: Formula — Mn0 2 . 

Manganese Dioxide is the common black ore of man- 
ganese, and is termed pyrolusite by mineralogists, it 
can be artificially formed by adding a solution of 
bleaching powder to a manganous salt. This substance 
yields one-third of its oxygen when heated to redness, 
forming the red oxide, and gives up half its oxygen 
when heated with sulphuric acid. It is largely used 
for the manufacture of chlorine. 

Manganese, Potassium and Oxygen Compounds. 

Manganic Acid: Formula — K 2 Mn0 4 . 
Permanganate of Potash: Formula — KMn0 4 . 

When an oxide of manganese is fused in the air with 
caustic alkali, a bright greeh mass is formed, which 
yields a dark green solution, this contains potassium 
manganate, or manganic acid, which may be crystal- 
lized, and is isomorphous with potassium sulphate and 



MERCURY 187 

chromate. If this green solution be allowed to stand, 
it slowly changes to a bright purple color, and hy- 
drated manganese dioxide is deposited, hence its com- 
mon name of mineral chameleon. It then contains a 
new salt in solution, a permanganate of potash, which 
may be obtained in the crystalline state by evapora- 
tion, and is isomorphous with potassium perchlorate. 
The presence of a few drops of acid at once effects this 
decomposition of the green solution. 

The manganates and permanganates readily give up 
a part of their oxygen in presence of organic matter, 
and they are largely used as disinfectants. 

MERCURY. 

Symbol— Hg. Atomic Weighi^-198.9. 

The principal ore of mercury is cinnabar, or Vermil- 
lion, which occurs in California, and also in China and 
Japan. The metal is easily obtained by roasting the 
ore, when the sulphur burns off as the dioxide, and the 
metal volatilizes, and its vapor is condensed in earthen 
pipes. Mercury is the only metal liquid at the ordinary 
temperature, it freezes at — 40°, in the solid state it 
is malleable. It boils at 350°, measured by the air 
thermometer, and gives off a slight amount of vapor 
at the ordinary temperature. Mercury when pure does 
not tarnish in moist or dry air, but when heated above 
300° it slowly absorbs oxygen, passing into the red 
oxide, and it combines directly with chlorine, bromine, 
iodine, and sulphur. Hydrochloric acid does not attack 
mercury, sulphuric acid, on heating, forms sulphur di- 
oxide and mercuric sulphate, nitric acid evolves nitric 
oxide, and forms mercuric nitrate. Mercury is largely 



188 ELEMENTARY CHEMISTRY 

used in the process of extracting gold and silver from 
their ores and in the arts, for silvering mirrors, and 
other purposes. Mercury is deposited from its solu- 
tions upon metallic iron or copper, in the form of a 
grey powder, which becomes bright on burnishing. 
Mercury salts act as valuable medicines. 

Mercury and. Chlorine Compounds. 
Mercuric Chloride: Formula — HgCl 2 . 

Mercuric Chloride, or Corrosive Sublimate, is pre- 
pared on a large scale by heating an intimate mixture 
of equal parts of mercuric sulphate and common salt. 
It is also formed when mercury burns in chlorine. It 
acts as a violent poison, it is soluble in water, fuses at 
265°, and boils at 295°. When ammonia is added to a 
solution of mercuric chloride, the so-called white pre- 
cipitate is thrown down. It is a chloride of mercury- 
ammonium. 

Mercurous Chloride or Calomel: Formula — Hg 2 Cl 2 . 

Mercurous Chloride or Calomel is generally prepared 
by heating a mixture of three parts of finely-divided 
metallic mercury with four parts of corrosive subli- 
mate, the metal combines with half the chlorine of the 
corrosive sublimate, and one atom of calomel is formed. 
The calomel sublimes, and is deposited in a solid cake. 
It must be finely ground and well washed, in order to 
free it from any soluble mercuric chloride which may 
remain undecomposed. Calomel is a white powder, in- 
soluble in water, it is decomposed by potash or am- 
monia. It is used largely in medicine. 



MOLYBDENUM 189 

Mercury and Oxygen Compounds. 
Mercuric Oxide: Formula — HgO. 

Mercury Monoxide, or Mercuric Oxide, is obtained 
by moderately heating the nitrate, or by heating the 
metal in the air for some time at a temperature of 
300°. The oxide thus prepared appears as a dark-red 
crystalline powder, by precipitating it from a solution 
of the nitrate by caustic potash, it falls as an amor- 
phous yellow powder. 

Mercury and Sulphur Compounds. 
Mercuric Sulphide or Cinnabar : Formula— HgS. 

Cinnabar or vermillion occurs native, and may be 
prepared artificially by heating a mixture of sulphur 
and mercury. When precipitated from a solution of a 
mercuric salt by sulphuretted hydrogen, the sulphide 
falls as a black amorphous powder but on sublimation 
it becomes red and crystalline. 

MOLYBDENUM. 

Symbol — Mo. Atomic Weight — 95.3. 

The chief ore of this metal is molybdenum disul- 
phide, a mineral in appearance resembling graphite. 
The metal is a grey substance, which oxidizes on 
heating in the air to molybdenum trioxide, a yellow 
powder which acts as an acid, forming with bases salts 
called molybdates. The compounds of molybdenum 
do not occur frequently, and are not used in the arts. 
Molybdic acid is, however, used as a reagent in the 
laboratory for detecting small traces of phosphoric 
acid. 



190 ELEMENTARY CHEMISTRY 

NICKEL. 
Symbol — Ni. Atomic Weight — 58.4. 

Nickel occurs in large quantities, combined with 
arsenic, as kupfernickel, also together with cobalt in 
speiss, and it is now prepared in considerable quanti- 
ties for the manufacture of German silver, an alloy of 
nickel, zinc, and copper. Nickel is a white, malleable, 
and tenacious metal, it melts at a somewhat lower 
temperature than iron, and is strongly magnetic, but 
loses this property when heated to 350°. 

Nickel and Oxygen Compounds. 

There are two oxides of nickel, the monoxide, NiO, 
and the sesquioxide, Ni 2 3 . The former of these gives 
rise to the nickel salts, which possess a peculiar apple- 
green color. The monoxide is obtained by heating the 
nitrate or carbonate, or by precipitating a soluble 
nickel salt with caustic potash, and heating the apple- 
green hydrate, NiH 2 2 , which is thrown down. The 
sesquioxide is a black powder, prepared by adding a 
solution of bleaching-powder to a soluble nickel salt. 

NITROGEN. 
Symbol— N. Atomic Weight— 13.9. Density— 13.9. 

Nitrogen is found in a free state in the atmosphere 
of which it forms about four-fifths by bulk, and occurs 
also in combination in many bodies such as ammonia, 
in the nitrates, and in many organic substances which 
form part of the bodies of vegetables and animals. 

Pure nitrogen cannot readily be prepared from air, 



NITROGEN 191 

as its separation from the gases of the helium group 
can only be effected by bringing the whole of the ni- 
trogen into the combined state. It can, however, be 
obtained from nitrogen compounds in a variety of 
ways. 

When a concentrated solution of ammonium nitrite 
is heated, the following reaction takes place: 
NH 4 N0 2 =N 2 +2H 2 0. 

It is more convenient to employ a mixture of sodium 
nitrite with ammonium sulphate, which by double de- 
composition yields sodium sulphate and ammonium 
nitrite, the latter then decomposing into nitrogen and 
water. The addition of a little potassium bichromate 
prevents the formation of any nitric oxide, and the 
best results are obtained with a solution of 1 part of 
sodium nitrite, 1-2 part of ammonium sulphate, and 
1 part of potassium bichromate, the gas being then 
washed through dilute sulphuric acid. 

Pure nitrogen is also formed by the action of chlorine 
upon ammonia; thus 

8NH 3 +3C1 2 =N 2 +6NH 4 C1. 

The chlorine evolved in a large flask passes into a 
three-necked Woulffe's bottle containing a strong 
aqueous solution of ammonia. The nitrogen gas which 
is here liberated is collected in the ordinary way over 
water, as shown in Fig. 26. Care must, however, be 
taken in this preparation that the ammonia is always 
present in excess, otherwise chloride of nitrogen may 
be formed, and this is a highly dangerous body, which 
explodes most violently. 

Nitrogen is also evolved in many other reactions 
and decompositions, such as the action of sulphuric 



192 ELEMENTARY CHEMISTRY 

acid on. a mixture of ammonium nitrate and glycerol 
and the decomposition by heating of ammonium bi- 
chromate. 

The oxygen of the air can readily be removed, yield- 
ing the mixture of gases formerly known as atmos- 
pheric nitrogen. 

A small light porcelain basin is allowed to swim 



Fig. 26. 

on the water of a pneumatic trough, a small piece of 
phosphorus brought into the basin and ignited, and 
the basin then covered by a large tubulated bell- jar 
as in Fig. 27. The phosphorus burns with the deposi- 
tion of a white cloud of phosphorus pentoxide, which, 
however, soon dissolves, whilst on cooling, one-fifth of 
the contents of the bell-jar is found to be filled with 
water. The colorless residual gas is nitrogen, this 



NITROGEN 



193 



may be easily proved by first equalizing the level of 
the water inside and outside the bell- jar, after which, 
on opening the stopper and plunging a burning taper 
into the bell-jar, the flame is seen to be instantly ex- 
tinguished. The nitrogen thus obtained is never per- 
fectly pure, as, in addition to elements of the helium 
group, it always contains small quantities of oxygen 
which have not been removed by the combustion of 
the phosphorus. In presence of aqueous vapor, phos- 




Fig. 27. 



phorus slowly absorbed the oxygen of the air, at tem- 
peratures above 15°. 

Pure nitrogen is a colorless, tasteless, inodorous gas, 
which is distinguished by its inactive properties, hence 
it is somewhat difficult to ascertain its presence in 
small quantities. As has been said, it does not sup- 
port combustion, nor does it burn nor render lime-water 
turbid. It combines directly with but very few non- 
metals, although indirectly it can easily be made to 
form compounds with most of these elements, and 
many of its compounds, such as nitric acid, ammonia 



194 ELEMENTARY CHEMISTRY 

and chloride of nitrogen, possess characteristic and 
remarkable properties. It combines directly with a 
number of the metals, such as lithium, calcium, barium, 
and magnesium, yielding compounds termed nitrides. 

Nitrogen forms an essential constituent of a very 
large number of animal and vegetable substances and 
is necessary for the maintenance of animal and vege- 
table life. It was however for a long time believed 
that members of the vegetable kingdom were unable 
to take up the free nitrogen of the air, and that they 
were dependent for their supply of this element on 
combined nitrogen contained in the atmosphere and 
in the soil chiefly in the form of nitric acid and am- 
monia. Certain leguminous plants, such as the white 
lupine, when grown in air free from ammonia and 
other nitrogen compounds, contain more nitrogen than 
was originally present in the seed and in the soil in 
which they were sown, and they must, therefore, have 
obtained the excess from the nitrogen of the air, since 
then it has been shown that certain algae, fungi, and 
mosses behave in a similar manner, and it is not im- 
probable that further investigation will show that this 
property is possessed by many other kinds of plants. 
This assimilation by plants is brought about by the 
action of certain micro-organisms, termed bacteroids, 
which are found in nodules on the roots, and absorb 
the free nitrogen from the air, forming compounds of 
nitrogen which are then assimilated by the plant. Other 
classes of micro-organisms also play a considerable 
part in the assimilation of nitrogen by plants, as many 
of the latter are incapable of directly assimilating am- 
monia, although they can take up nitric acid, the mi- 
cro-organisms in the soil bring about the conversion 



NITROGEN 195 

of ammonia into nitric acid ; one species of these con- 
verting the former into nitrous acid, and another con- 
verting the nitrous into nitric acid. 



Nitrogen and Chlorine Compounds. 
Nitrogen Chloride: Formula — NCl r 

Chloride of nitrogen is a thin yellowish oil which 
evaporates quickly on exposure to the air and possesses 
a peculiar smell, the vapor attacking the eyes and 
mucous membrane violently. 

The following method is employed for preparing 
this dangerous substance in small quantities, and for 
showing its explosive properties without risk. A flask 
of about two litres capacity, having a long neck, is 
filled with chlorine, and placed mouth downwards in 
a large glass basin filled with warm saturated solution 
of sal-ammoniac as shown in Fig. 28. Below the neck 
of the flask is placed a small thick leaden saucer, in 
which the nitrogen chloride is collected. The solution 
of sal-ammoniac absorbs the chlorine, and, as soon as 
the flask is three parts filled by the liquid, oily drops 
are seen to collect on the surface inside the flask. These 
gradually increase, and at last drop one by one down 
into the leaden saucer. When a few drops have col- 
lected, the leaden saucer may be carefully removed by 
a pair of clean tongs, another being placed in its stead. 
A small quantity of the chloride of nitrogen may be 
exploded by touching it w T ith a feather moistened with 
turpentine attached to the end of a long rod. An- 
other drop of the oil may be absorbed by filtering 
paper, and when this is held in a flame a loud explo- 
sion likewise ensues. 



196 



ELEMENTARY CHEMISTRY 



The force of the explosion may be rendered still 
more evident by placing the flask in a strongly con- 
structed box with glass sides, and when the drops of 
nitrogen chloride collect on the surface of the solution, 
passing in some turpentine from a stoppered funnel, 
the latter is kept outside the box and is connected with 




Fig. 28. 

it by means of rubber and glass tube the end of which 
dips under the neck of the flask. When the turpen- 
tine reaches the surface of the solution in the flask a 
bright flash is seen, and a violent thunderlike explo- 
sion occurs, completely shattering the flask. 



Nitrogen and Hydrogen Compounds. 
Ammonia, NH 3 — Hydrazine, N 2 H 4 and Azoimide, N 3 H. 



NITROGEN 197 

Ammonia: Formula— NH r Molecular Weight — 16.93. 
Density— 8.45. 

Ammonia is chiefly obtained from the decomposition 
of animal or vegetable matters containing nitrogen and 
hydrogen, being formed either gradually at the or- 
dinary temperature, or quickly under the influence of 
heat. When horns, or clippings of hides, or coal is 
heated, ammonia is given off; hence ammonia was 
known as spirits of hartshorn. The name ammonia is 
derived from the fact that a compound containing am- 
monia, called sal-ammoniac, was first prepared by the 
Arabs in the deserts near the temple of Jupiter Ammon, 
by heating camels' dung. Guano, the dried excrement 
of sea-birds, and the urine of animals, likewise contain 
large quantities of ammonia. Ammonia and its com- 
pounds are now, however, mainly obtained from the 
ammoniacal liquors of gasworks. Coal contains about 
2 per cent of nitrogen, which, when the coal is heated 
in close vessels, mostly comes off in combination with 
the hydrogen of the coal as ammonia. Hydrochloric 
acid is added to this ammoniacal liquor, and the solu- 
tion evaporated, when the sal-ammoniac of commerce 
is obtained. 

Ammonia may also be formed by the action of nas- 
cent hydrogen on dilute nitric acid, and when this acid 
is placed in contact with metallic zinc or iron, am- 
monia is formed. 

Ammonia gas is best prepared by heating in a glass 
flask one part of sal-ammoniac, or ammonia hydro- 
chlorate, and two parts of powdered quicklime. 

Quicklime and sal-ammoniac give calcium chloride, 
and ammonia and water. 



198 ELEMENTARY CHEMISTRY 

Ammoniacal gas is colorless, and possesses a most 
pungent and peculiar smell, by means of which it can 
be readily recognized, it is lighter than air, and it may 
be collected by displacement, the neck of the bottle in- 
tended to receive the gas being turned downwards, as 
in Fig. 29. A cylinder filled with quicklime is here 







Fig. 29. 

placed between the flask and the bottle, for the purpose 
of completely drying the ammonia. Ammonia may 
also be collected over mercury, but not over water, as 
it is extremely soluble in this liquid, one gram of water 
at 0° absorbing 0.877 gram, or 1149 times its volume, 
of ammonia, under a pressure of 760 mm., whilst at 
20° the same weight of water absorbs 0.520 gram, or 



NITROGEN 199 

681.1 times its volume, under the same pressure. The 
solution of ammonia gas in water is the common liquor 
ammoniac of commerce. Ammonia gas, as well as the 
aqueous solution, possesses a strong alkaline reaction, 
turning red vegetable colors blue. It unites with the 
most powerful acids, forming compounds called the 
salts of ammonia, which closely resemble the salts of 
the alkaline metals, hence the name of the volatile al- 
kali has been given to ammonia. 

On exposure to a pressure of seven atmospheres at 
the ordinary temperature of the air, about 15° Centi- 
grade, ammonia condenses to a colorless liquid, boiling 
at — 38.5°, and this liquid, if cooled below — 75°, 
freezes to a transparent solid. An elegant application 
of the principle of the latent heat of vapors has re- 
cently been made in the case of ammonia in M. Carre's 
freezing machine. This consists essentially of two 
strong iron vessels connected in a perfectly air-tight 
manner by a bent pipe, one of these vessels contains 
an aqueous solution of ammonia saturated with the 
gas at 0°. When it is desired to procure ice, the ves- 
sel containing the ammonia solution is gradually heated 
over a large gas burner, the other vessel being placed 
in a bucket of cold water. In consequence of the in- 
crease of temperature, the gas cannot remain dissolved 
in the water, and passes into the receiver, where, as 
soon as the pressure amounts to about 10 atmospheres, 
it condenses in the liquid form. When the greater part 
of the gas has thus been driven out of the water, the 
apparatus is reversed, the retort being cooled in a cur- 
rent of cold water, whilst the liquid it is desired to 
freeze is placed in the interior of the receiver. A re- 
absorption of the ammonia by the water now takes 



200 ELEMENTARY CHEMISTRY 

place, and a consequent evaporation of the liquefied 
ammonia in the receiver. This evaporation is accom- 
panied by an absorption of heat which becomes latent 
in the gas, hence the receiver is soon cooled far below 
the freezing point, and ice is produced around it. 

The composition of ammonia may be ascertained by 
leading the gas through a red-hot tube, or passing a 
series of electric sparks through the gas, when it will 
be decomposed into nitrogen and hydrogen, which will 
be found to occupy together a volume twice as large as 
the ammonia taken, and to be mixed together in the 
proportions of three volumes of hydrogen to one vol- 
ume of nitrogen. 

Hydrazine or Diamide: Formula — N 2 H 4 . 

Free hydrazine is difficult to prepare on account of 
the great stability of the hydrate, which is formed 
whenever the base is liberated in presence of water. 

It is, however, best prepared by gradually adding 
the hydrate to anhydrous barium oxide, heating for 
some time at 110 — 120° and then distilling under 150 
— 100 mm. pressure, the air of the apparatus being dis- 
placed by hydrogen so as to prevent the oxidation of 
the base. The distillate is not quite free from water 
and must therefore be again distilled over baryta. The 
free base is a colorless liquid, which boils at 113.5° 
and readily solidifies forming crystals which melt at 
1.4°. The base is extraordinarily hygroscopic and 
mixes readily with water or alcohol, but is only very 
sparingly soluble in organic solvents generally. It 
fumes in the air and is readily inflammable, but not 
explosive. It acts as a most violent reducing agent, 
and bursts into flame when brought into contact with 



NITROGEN 201 

chlorine, free nitrogen and hydrochloric acid being 
formed, on bromine and iodine it also reacts vigor- 
ously, forming the corresponding acids. It slowly 
oxidizes in the air and when strongly heated decom- 
poses, the final products being ammonia and nitrogen, 
3N 2 H 4 =N 2 +4NH 3 . 

Free hydrazine dissolves many salts and yields solu- 
tions which conduet electricity, so that its action in 
this respect is comparable with that of water and am- 
monia. 

Azoimide or Hydrazoic Acid: Formula — N.H. 

A convenient source of azoimide is the crude amido- 
guanidine obtained in the preparation of hydrazine, 
this substance is converted by nitrous acid into diazo- 
guanidine nitrate, which on boiling with alkalies yields 
azoimide and cyanamide, the former may be isolated 
by acidification and distillation. 

A further very interesting synthesis of azoimide from 
purely inorganic sources is by obtaining its sodium salt 
by the action of nitrous oxide on sodamide. 

The water formed acts on a further molecule of soda- 
mide yielding caustic soda and ammonia. The soda- 
mide is obtained by passing ammonia over metallic 
sodium at a temperature of 150 — 250°, ai^d as soon as 
all metallic sodium has disappeared, the steam of am- 
monia is replaced by one of nitrous oxide and contin- 
ued till ammonia is no longer evolved, the product is 
then dissolved in water and distilled with dilute sul- 
phuric acid. 

Several other interesting modes of formation have 
also been discovered. A yield of 36 per cent, of the 



202 ELEMENTARY CHEMISTRY 

theoretical amount is obtained when a benzene solution 
of nitrogen chloride is shaken with a solution of hydra- 
zine sulphate, and caustic soda added at intervals. It 
is also formed when a mixture of molecular proportions 
of hydrazine and hydroxylamine dissolved in dilute 
sulphuric acid is oxidized by chromic acid or hydrogen 
dioxide. The silver salt may also be conveniently pre- 
pared by cautiously warming hydrazine sulphate with 
nitric acid and passing the gas evolved into silver ni- 
trate solution. 

Nitrogen and Oxygen. 
The Air. 

That the oxygen and nitrogen of the air are mechani- 
cally mixed and not chemically combined is seen from 
the following facts: 

The quantities of nitrogen and oxygen in the air do 
not present any simple relation to the atomic weights 
of these elements, and the proportions in which they 
are mixed are variable. 

On mixing oxygen and nitrogen gases mechanically 
in the proportion in which they occur in air, no con- 
traction or evolution of heat is observed, and the mix- 
ture behaves in every way like air. 

When air is dissolved in water, the proportion be- 
tween the oxygen and nitrogen in the dissolved air is 
quite different from that of the undissolved air, the 
difference being in strict accordance with the laws of 
gas-absorption on the assumption that the air is a mix- 
ture. When water is saturated with air at any tem- 
perature below 30°, the following is the proportion of 
oxygen and nitrogen contained in the dissolved and the 
original air: 



NITROGEN 



203 





Air dissolved 
in water. 


Air undissolved ft 
in water. 


Oxygen 
Nitrogen 


35.1 

64.9 

100.0 


20.96 
79.04 

100.00 



If the air were a chemical combination of oxygen and 
nitrogen, such a separation by solution would be im- 
possible, 




Fig". 30. 

When liquefied air is allowed to boil, the nitrogen 
passes off much more rapidly than the oxygen, which 
could not take place if the two gases were chemically 
combined. 

In order to show the composition of the air, the ap- 
paratus Fig. 30 is used. This consists of a calibrated 
and divided glass tube filled to a given point with air 
over mercury. Into this is introduced a small piece of 
phosphorus supported upon a copper wire. Gradually 
all the oxygen is absorbed and the mercury rises in the 



204 



ELEMENTARY CHEMISTRY 



tube. After a while the volume of residual gas is read 
off, and, corrections having been made for temperature 
and pressure, it is found that 100 volumes of the air 
contain about 21 volumes of oxygen. 

Another less exact but more rapid method of ex- 
hibiting the same fact is carried out by help of the 




Fig. 31. 

arrangement shown in Fig. 31. In the beaker- glass c 
is placed the iron stand d carrying the iron cup e, con- 
taining a small piece of phosphorus. Over this stand 
is placed the cylinder a. The upper part of this cylin- 
der is graduated into five equal divisions, and water 
is poured into the beaker- glass until the level reaches 
the first division. The phosphorus in the cup is ignited 






NITROGEN 205 

by dropping down on it a chain which has been heated 
in a flame. The phosphorus then burns, the fumes of 
phosphorus pentoxide are absorbed by the water, and, 
when the gas has cooled, and the pressure been equal- 
ized by bringing the level of the water outside up to 
that inside the cylinder, four-fifths of the original vol- 
ume of the air remain unabsorbed. 

Analyses of air collected in various parts of the 
globe, made with the greatest care have shown that 
the relative quantities of oxygen and nitrogen remain 
the same, or very nearly the same, from whatever re- 
gion the air may have been taken. So that whether the 
air be deri ved from the tropics or the arctic seas, from 
the bottom of the deepest mine or from an elevation of 
20,000 feet above the earth's surface, it contains from 
20.9 to 21 volumes of oxygen per cent. 

When we know the composition of air by volume, and 
the relative densities of the two constituent gases, 14 
for nitrogen and 16 (actually 15.88, when hydrogen is 
1), for oxygen, we can calculate its composition by 
weight, we thus find that in 100 grams of air, 23.16 
grams of oxygen are mixed with 76.84 grams of nitro- 
gen. It is important to control this calculation by 
experiment, for this purpose a large glass globe fur- 
nished with a stopcock is rendered vacuous by the air- 
pump and then weighed, a tube of hard glass filled with 
copper turnings and also furnished with stopcocks is 
likewise weighed. This tube is then heated to redness 
in a long tube-furnace, and connected at one end with 
the empty flask, at the other with a series of tubes filled 
with caustic potash and sulphuric acid, for the pur- 
pose of completely freeing the air passing through 
them from carbonic acid and aqueous vapor, the cocks 



206 ELEMENTARY CHEMISTRY 

are then slightly opened, and air allowed to pass slow- 
ly through the purifiers into the hot tube, where it is 
completely deprived of oxygen by the hot metallic 
copper, which is thereby oxidized, the nitrogen passing 
on alone into the empty flask. After the experiment 
is concluded, the cooled tube is again weighed, and the 
increase over the former weighing gives the quantity 
of oxygen, whilst the increase in weight of the globe 
gives the nitrogen. The mean of a large number of 
experiments thus made shows that 100 parts by weight 
of air contained 23 parts by weight of oxygen and 
77 of nitrogen. 

In addition to the two above-mentioned gases, the 
air contains several other important constituents, es- 
pecially carbonic acid gas, aqueous vapor, and am- 
monia gas. The carbonic acid gas of the air plays an 
important part in the phenomena of vegetation, this 
gas being the source from which plants obtain the car- 
bon they need to form their tissues. The quantity of 
carbonic acid present in the air is very small compared 
with the quantities of oxygen and nitrogen, being only 
about 4 volumes to 10 ; 000 of air, nevertheless the ab- 
solute quantity of this gas contained in the whole at- 
mosphere is enormously large, about 3,000 billion kilos. 
The quantity of carbonic acid contained in the air can 
be found by drawing a known volume of perfectly dry 
air, not less than 20 litres, through weighed tubes con- 
taining caustic potash, the increase in weight of the 
tubes gives the weight of carbonic acid contained in 
the air drawn through. Hail is caused by the congela- 
tion of raindrops in passing through a stratum of air 
below the freezing point. The quantity of rain thus 
deposited is very large : 1 cubic metre of air saturated 



NITROGEN 207 

with moisture at 25° Centigrade contains 22.5 grams 
of water, and if the temperature of this air be reduced 
to 0° Centigrade, it will then be capable of retaining 
only 5.4 grams of water vapor, hence 17.1 grams of 
water will be deposited as rain. Instruments for as- 
certaining the degree of moisture or humidity of the 
air are termed hygrometers. 

The deposition of dew is caused by the rapid cooling 
of the earth's surface by radiation after sunset, and by 
the consequent cooling of the air near the ground be- 
low the temperature at which it begins to deposit mois- 
ture. 

The amount of aqueous vapor contained in the air 
at any time can be determined by the apparatus used 
for the estimation of the carbonic acid, for the mois- 
ture must be removed from the air before the carbonic 
acid can be absorbed, and the increase in weight of 
the tubes filled with pumice-stone moistened with sul- 
phuric acid gives the weight of aqueous vapor. In 
general the air contains from 50 to 70 per cent of the 
quantity necessary to saturate it. If the quantity be 
not within these limits, the air is either unpleasantly 
dry or moist. 

The next important constituent of the air is ammo- 
nia, which is a compound of nitrogen and hydrogen, 
and only exists in comparatively very minute quanti- 
ties, about 1 part in 1,000,000 of air. Nevertheless it 
plays a very important part, as it is mainly from this 
ammonia that vegetables obtain the nitrogen which 
they need to form their seeds and fruit, for it appears 
that all plants have not the power of assimilating the 
free nitrogen of the atmosphere. Other substances 
which occur in the atmosphere in very small quantities 



208 ELEMENTARY CHEMISTRY 

may be considered as accidental impurities. Amongst 
them, volatile organic matter is the most important, as 
probably influencing to a great extent the healthiness 
of the special situation. We become aware of the exist- 
ence of such organic putrescent substances when en- 
tering a crowded room from the fresh air, and it is 
probable that the well-known unhealthiness of marshy 
and other districts is owing to the presence of some 
organic impurity. At present, however, we possess but 
little certain knowledge on this subject. Ozone is also 
present in fresh air, but generally absent in the close 
air of towns and dwelling-rooms, owing to its decom- 
position by the organic matter in such air, we do not 
know how it is formed in nature, unless it be by the 
discharge of atmospheric electricity. 

Nitrogen and Oxygen Compounds. 

Nitrogen Monoxide or Nitrous Oxide — N 2 0, 
Nitrogen Dioxide or Nitric Oxide — NO, 
Nitrogen Trioxide — N 2 3 , 
Nitrogen Peroxide — N0 2> 
Nitric Anhydride — N 2 5 . 

Nitrous Oxide: Formula — N 2 0. Molecular Weight — 
43.7. Density— 21.85. 

Nitrous Oxide is called laughing gas, because when 
it is mixed with air and inhaled, it produces a peculiar 
and transient intoxicating effect, and when inhaled in 
its pure state for a few minutes, it acts as an anaes- 
thetic, and renders the person insensible to pain for a 
short time. 

Nitrate of ammonia which is prepared by neutral- 



NITROGEN 209 

izing nitric acid with ammonia, gives off nitrous oxide 
when decomposed by heat as in Fig. 32, thus 
HN0 3 +NH 3 =(NH 4 )N0 3 

and (NH 4 )X03=N 2 0+2H 2 °0. 
That is, ammonium nitrate yields nitrous oxide and 
water. 

Nitrous oxide is a colorless inodorous gas possessing 
a slightly sweet taste, it is somewhat soluble in cold 
water, one volume of water at 0° dissolving 1.305 




Fig. 32. 

volumes of the gas, whilst one volume of water at 24° 
dissolves only 0.608 volume. Nitrogen monoxide dif- 
fers from all such gases which we have previously con- 
sidered, inasmuch as it liquefies when exposed either to 
great pressure or to an intense degree of cold. Thus, 
if it be brought under a pressure of about 30 atmos- 
pheres at 0°, or if it be cooled down to — 88° under the 
ordinary pressure, it forms a colorless liquid, in other 
words, the pressure of nitrous oxide vapor or gas is 1 
atmosphere at — 88°, and 30 atmospheres at 0° Centi- 



210 ELEMENTARY CHEMISTRY 

grade. If this liquid be cooled below — 115°, it solidi- 
fies to a transparent mass. By the rapid evaporation 
of this liquid in vacuo, the lowest artificial temperature 
hitherto known has been attained, about — 140° Centi- 
grade. 

A glowing chip of wood when plunged into nitrous 
oxide rekindles, and the wood continues to burn with 
a brighter flame than in the air, whilst phosphorus on 
burning in this gas evolves nearly as much light as in 
pure oxygen, a feeble flame of sulphur is, however, ex- 
tinguished on bringing it into this gas, but if burning 
strongly it also continues to burn brightly. This is ow- 
ing to the fact that the gas has to be decomposed into 
nitrogen, one volume, and oxygen, half a volume, be- 
fore bodies can burn in it, and to effect this decompo- 
sition a tolerably high temperature is necessary. 

Nitric Oxide : Formula — NO. Molecular Weight — 
29.8. Density— 14.9. 

A colorless gas obtained by acting upon copper turn- 
ings with nitric acid. Copper and nitric acid give cop- 
per nitrate, nitrogen dioxide, and water. 

This substance has not been condensed to a liquid, in 
contact with oxygen it combines directly with this gas, 
forming red fumes which are readily soluble in water, 
and by this property it may be distinguished from all 
other gases. Although nitric oxide contains half its 
volume of oxygen, and more oxygen in proportion by 
weight than nitrous oxide, it does not easily support 
combustion, as it requires a high temperature for its 
decomposition, thus, ignited phosphorus, unless burning 
very brightly, is, extinguished on plunging it into nitric 
oxide gas. 



NITROGEN 211 

The composition of this gas may be determined as 
follows: One volume of nitrogen dioxide yields half 
a volume of nitrogen, as the weight of one volume of 
nitrogen dioxide is 15, the weight of oxygen contained 
in one volume of this gas is 15 — 7=8, or two volumes 
of nitrogen dioxide weigh 30, and are composed of one 
volume of nitrogen weighing 14, and one of oxygen 
weighing 16 (actually 15.88 when hydrogen is 1). 
Hence, in accordance with the law respecting the densi- 
ties of compound gases, the formula of this oxide 
should be NO and not N 2 2 , the physical properties of 
the gas likewise, compared with those of nitrous oxide, 
seem to indicate that this latter has a more compli- 
cated constitution. 

If N 2 2 is used for the formula then this gas will be 
an exception to the law that the density of any com- 
pound gas is half its molecular weight. The density 
being 15, the molecular weight, according to this law, 
is 30 and the formula is NO. It is as though the single 
molecule N 2 2 had split up into two molecules of NO. 
The name nitrogen dioxide is given to this gas because 
for the same weight of nitrogen it contains twice as 
much oxygen as nitrogen monoxide. 

Nitrogen Trioxide: Formula — N 2 3 . Molecular 
Weight— 75.5, Density— 37.7. 

This substance is prepared by mixing four volumes 
of dry nitrogen dioxide with one volume of oxygen, 
and cooling the mixture to — 18°, the two gases com- 
bine to form red fumes, which condense to a volatile 
indigo-blue colored liquid, the same blue body is ob- 
tained by adding water to nitrogen peroxide and dry- 
ing the distillate over calcium chloride. It is also 



212 ELEMENTARY CHEMISTRY 

formed by the action of moderately strong nitric acid 
upon arsenic trioxide, with formation of arsenic acid. 
Arsenic trioxide and nitric acid and water yield nitro- 
gen trioxide and arsenic acid. 

Nitrogen trioxide dissolves in ice-cold water, form- 
ing a blue liquid, and containing nitrous acid or hydro- 
gen nitrite, HN0 2 , in solution, this compound is very 
unstable, and decomposes when the water is warmed, 
into nitric acid and nitric oxide. 

The salts formed by nitrous acid are, however, not 
liable to such easy decomposition, potassium nitrite, 
KN0 2 , is obtained by heating potassium nitrate, KN0 3 , 
which loses one atom of oxygen, the same salt is pro- 
duced when nitrogen trioxide is led into a solution of 
caustic potash. 

Hence, nitrogen trioxide stands to the nitrites in the 
same position as nitrogen pentoxide to the nitrates. It 
will be noticed that nitric acid forms salts called ni- 
trates, whilst nitrous acid gives rise to nitrites, this is 
an example of a general rule adopted in chemical 
nomenclature that if the specific name of an acid or hy- 
drogen salt end in ous, the names of the corresponding 
metallic salts end in ite, whilst acids whose names end 
in ic form salts ending in ate. 



Nitrogen Peroxide: Formula — N0 2 . Molecular 
Weight— 45.7. Density— 22.8. 

The red fumes which are formed when nitric oxide 
comes into contact with oxygen or air, consist chiefly 
of nitrogen peroxide. If one volume of dry oxygen be 
mixed with two volumes of dry nitric oxide and the 
red fumes which are produced led into a tube sur- 



NITROGEN 



213 



rounded by a freezing mixture, the peroxide condenses 
in the tube .either as a liquid or in the form of crystals. 

Nitrogen peroxide is also formed by the decomposi- 
tion which many nitrates undergo when heated, and is 
usually prepared by strongly heating lead nitrate in 
a retort of hard glass as shown in Fig. 33. 

This mode of preparation is, however, not very con- 
venient, and a considerable loss of material occurs, as 
the oxygen which is evolved carries away some quan^ 




Fig. 33. 

tity of the peroxide even when the tube into which the 
fumes are led is plunged into a freezing mixture. 

The following method is free from the above objec- 
tions. Arsenious oxide (white arsenic) in the form of 
small lumps is placed in a flask and covered with ordi- 
nary nitric acid, or, with a mixture of nitric acid of 
specific gravity 1.5 and half its weight of sulphuric 
acid, the red fumes, which are given off in quantity on 
gently heating, are led into a receiver surrounded by 



214 ELEMENTARY CHEMISTRY 

a freezing mixture, where a mixture of trioxide and 
tetroxide of nitrogen collects. The mixture is freed 
from the trioxide by the addition of strong nitric acid 
and a large quantity of phosphorus pentoxide^ and the 
tetroxide is then poured off from the syrupy layer, and 
distilled. The trioxide may also be converted into per- 
oxide by a current of air. 

Nitrogen peroxide may also be prepared by mixing 
the fumes obtained by the action of nitric acid of spe- 
cific gravity 1.4 on arsenious oxide with oxygen, dry- 
ing over calcium nitrate, condensing and redistilling. 

Nitric Anhydride: Formula — N 2 05. 

This compound is a white crystalline solid obtained 
by removing the elements of water from nitric acid by 
means of phosphorus pentoxide, a substance which has 
a great power of abstracting Water. 

Nitrogen pentoxide, or nitric anhydride, is an unsta- 
ble body, and is not used in the arts, it is interesting 
as being the highest oxide of nitrogen, and the anhy- 
dride of nitric acid, uniting with great energy with 
water to form this acid. 

OXYGEN. 

Symbol — 0. Atomic Weight — 15.88. Density — 15.9. 

Oxygen is a colorless invisible gas, possessing neither 
taste nor smell. It exists in the free state in the at- 
mosphere, of which it constitutes about one-fifth by 
bulk, whilst, in combination with the other elements, 
it forms nearly half the weight of the solid earth, and 
eight-ninths by weight of water. The birth of the 
modern science of chemistry may be dated from the 



OXYGEN 



215 



discovery of oxygen. Oxygen gas can be prepared 
from the air, but it is more easily obtained from many 
compounds which contain it in large quantities. Oxy- 
gen may be prepared by heating red mercury oxide. 
This substance is made up of 200 parts by weight of 
mercury, and sixteen parts of oxygen, when strongly 
heated, it is decomposed, and yields metallic mercury 
and oxygen gas. Oxygen can be more cheaply ob- 
tained by heating potassium chlorate, a white salt 




Fig. 34. 

which yields on heating 39.2 per cent of its weight of 
this gas. In order to collect the oxygen thus given 
off, powdered potassium chlorate is placed in a small 
thin glass flask, furnished with a well-fitting cork, into 
which a bent tube is inserted. The lower end of the 
tube dips under the surface of water in a pneumatic 
trough, and the gas, on being evolved, bubbles out 
from the end of the tube, and is collected in jars or 
bottles filled with water, and placed with their mouths 
downwards in the trough. Fig. 34 shows the arrange- 
ment of the apparatus needed for the preparation of 



216 ELEMENTARY CHEMISTRY 

oxygen gas. If a small quantity of manganese dioxide 
(black oxide of manganese) be mixed with the potas- 
sium chlorate, the oxygen is given off from the chlo- 
rate at a much lower temperature, and the evolution 
of the gas is facilitated, but the manganese dioxide 
undergoes no change whatever. 

All the elements, with the single exception of fluo- 
rine, combine with oxygen to form oxides. In this act 
of combination, which is termed oxidation, heat is al- 
ways, and light is frequently, given off. When bodies 
unite with oxygen,, evolving light and heat, they are 
said to burn, or undergo combustion. All bodies which 
burn in the air burn with increased brilliancy in oxy- 
gen gas, and many substances, such as iron, which -do 
not burn in the air, may be made to do so in oxygen. 
A redhot chip of wood, or a taper with glowing wick, 
is suddenly rekindled and bursts into flame when 
plunged into a jar of this gas. Sulphur, which in the 
air burns with a pale lambent flame, emits in oxygen 
a bright violet light, and a small piece of phosphorus, 
w T hen inflamed and placed in oxygen, burns with a daz- 
zling light. If the jars in which these experiments 
have been performed be afterwards examined, it is 
found that the substances produced by combustion in 
oxygen possess acid characters, they have the power 
of turning red certain vegetable blue coloring matters, 
such as litmus. A bundle of fine iron wire can be easily 
burnt in oxygen by tipping the end with burning sul- 
phur, and then plunging the iron thus tipped into a 
jar of the gas, the oxide of iron, formed by the com- 
bustion, drops down in the molten state. 

Many other substances may be employed for the 
preparation of oxygen, if large quantities of the gas 



OXYGEN 217 

are needed, manganese dioxide may be heated to red- 
ness in an iron bottle, 100 parts by weight of the oxide 
yield 12.3 by weight of oxygen. Another interesting 
decomposition by which oxygen is set free is that ef- 
fected by sunlight upon the carbonic acid gas contained 
in the air, this is accomplished by means of the green 
coloring matter of plants. Sunlight has the power, in 
presence of this green coloring matter, of decomposing 
carbonic acid, the carbon is taken up by the plant for 
its growth, while the oxygen is set free, and is after- 
wards used by animals for the support of the process 
of respiration. In the act of inspiration, animals 
breathe in the oxygen of the air, whilst in that of ex- 
piration, they breathe out carbonic acid gas. Hence 
oxygen is necessary to animal life, this gas was for- 
merly termed vital air. The chemical change which 
oxygen effects upon the body of the animal is in fact 
identical with that which goes on when a piece of char- 
coal burns in the air or oxygen, this may be rendered 
evident by a simple experiment. If some clear lime- 
water be poured into a bottle of oxygen in which char- 
coal has been burnt, the lime-water will become milky, 
owing to the formation of a compound of lime and car- 
bonic acid (called chalk), this acid being produced by 
the combustion, if the air contained in the lungs be 
next blown through a piece of glass tubing into some 
more clear lime-water, a turbidity (from the formation 
of chalk) will at once occur, proving that carbonic 
acid gas is given off from the lungs. This carbonic 
acid arises from the oxidation of the constituents of 
the body, and by this oxidation the heat of the body, 
which is greater than that of surrounding inanimate 
objects, is sustained. When this chemical process stops 



218 ELEMENTARY CHEMISTRY 

the animal dies, and the temperature of the body sinks 
to that of the neighboring objects. Carbonic acid, ni- 
trogen, and some other gases cause death when inhaled, 
because they do not contain free oxygen, and hence 
the process of oxidation in the body ceases. This cause 
of death is independent of any poisonous action of the 
gases. 

OZONE. 

Symbol— O r Molecular Weight— 47.64. Density— 23.8. 

If a series of electric discharges be sent through a 
tube containing pure and dry oxygen, only a small 
portion of the gas is converted into ozone. If the 
ozone, however, is removed as soon as formed, by a so- 
lution of iodide of potassium, for example, the whole 
of the oxygen can be gradually converted into ozone. 
In order to obtain the maximum production of ozone, 
pure oxygen gas is allowed to pass through an ap- 
paratus, Fig. 35, which consists essentially of an iron 
tube BB turned very truly on the outside, through 
which a current of cold water can be passed by means 
of the tubes CC. Outside this metal cylinder is one of 
glass AA very slightly larger than the iron one. By 
means of tubes DD air or oxygen can be passed through 
the annular space between the two cylinders. Part of 
the outer cylinder at G is covered with tinfoil. The 
outer tinfoil coating and the inner metal cylinder are 
connected with the poles of an induction coil at E and 
F. By this means the oxygen is subjected to a series 
of silent discharges, by which it is converted partially 
into ozone. Oxygen gas through which an electric 
spark has been passed possesses a peculiar smell, and 
at once tarnishes a bright surface of mercury. This 



OZONE 



219 



peculiar strongly-smelling substance, which is called 
ozone, is capable of liberating iodine from potassium 
iodide, and of effecting many other oxidizing actions. 
Ozone is also produced in many other ways. 

It is evolved at the positive pole in the electrolysis of 
acidulated water. 

It is formed by the discharge from an electrical ma- 
chine through air or through oxygen gas. 

When fluorine is passed into water at 0°, the oxygen 
liberated contains 10 to 14 per cent of ozone. 




Fig. 35. 



By acting with strong sulphuric acid upon dry bari- 
um dioxide, oxygen is given off which contains a con- 
siderable quantity of ozone. 

When oxygen is passed over heated manganese diox- 
ide, cobalt oxide or certain other metallic oxides, ozone 
is produced. 

Ozone is obtained when potassium permanganate and 
sulphuric acid are distilled together in a vacuum. 

It is stated that ozone is formed during combustion, 



220 ELEMENTARY CHEMISTRY 

and can be recognized by its smell when a current of 
air is blown through the upper portion of a flame. 

Powerful oxidizing agents are produced on the oxi- 
dation of phosphorus, turpentine and other oils by at- 
mospheric oxygen, and these were formerly thought to 
be ozone. 

PLATINUM. 
Symbol— Pt. Atomic Weight— 193.4. 

Platinum is a comparatively rare metal, which al- 
ways occurs in the native state, and generally alloyed 
with other metals. This alloy occurs in small grains 
in detritus and gravel in Siberia and Brazil, it has not 
been found in situ in the original rock, which probably 
belongs to the old plutonic series. 

The original mode of obtaining the metal was to 
dissolve the ore in aqua regia, and precipitate the plati- 
num with sal-ammoniac, as the insoluble double chlo- 
ride of ammonium and platinum. This precipitate, on 
heating, yields metallic platinum in a finely divided or 
spongy state, and this sponge, if forcibly pressed and 
hammered when hot, gradually assumes a coherent 
metallic mass, the particles of platinum welding to- 
gether, when hot, like iron. A new mode of preparing 
the metal has recently been employed, the ore being 
melted in a very powerful furnace. In this way a pure 
alloy of platinum, iridium, and rhodium is formed, the 
other constituents and impurities of the ore either being 
volatilized by the intense heat, or absorbed by the lime 
of which the crucible is composed. This alloy is in 
many respects more useful than pure platinum, being 
harder and less easily attacked by acids than the pure 
metal. 



PLATINUM 221 

Platinum possesses a bright white color, and does 
not tarnish under any circumstances in the air, it is 
extremely infusible, and can only be melted by the 
heat of the oxy-hydrogen blowpipe. It is unacted upon 
by the ordinary acids, but dissolves in aqua regia, and 
hence platinum vessels are much used in the labora- 
tory. Caustic alkalies, however, act upon the metal at 
high temperatures. When finely divided, metallic 
platinum has the power of condensing gases on to its 
surface in a remarkable degree. 

Platinum and Oxygen Compounds. 

Platinum and oxygen unite in two proportions to 
form Platinum monoxide, PtO, and Platinum dioxide, 
Pt0 2 . The first of these oxides is a black powder, eas- 
ily decomposed on heating, and yielding a series of 
unstable salts, the second is obtained as a brown hy- 
drate, by adding to a solution of platinic nitrate half 
its equivalent of caustic potash, the hydrate, when 
heated, first loses its water, forming the anhydrous 
oxide, and then parts with its oxygen, leaving the 
metal platinum. 

Platinum, and Chlorine Compounds. 

Platinum dichloride: Formula — PtCl 2 . 

Platinum tetrachloride: Formula — PtCl 4 . 

Platinum dichloride is a green insoluble powder, ob- 
tained by heating the higher chloride to 200°. Plati- 
num tetrachloride is the most important platinum com- 
pound. It is obtained as a yellowish-red solution by 
dissolving the metal in aqua regia, on evaporation, 
crystals of a compound of platinum tetrachloride with 



222 ELEMENTARY CHEMISTRY 

hydrochloric acid separate out. Platinum tetrachlo- 
ride combines with many alkaline chlorides to form 
double salts. These compounds with potassium and 
ammonium are insoluble in water, and are isomorphous, 
crystallizing in cubes, while the sodium salt is soluble, 
and crystallizes in large prisms. 

POTASSIUM. 
Symbol— K. Atomic Weight— 38.85. 

The metal potassium was discovered by the decom- 
position of the alkali potash into potassium, hydrogen, 
and oxygen, by means of a powerful galvanic current. 
Before this time the alkalies and alkaline earths were 
supposed to be elementary bodies. The metal is now 
prepared by heating together potash and carbon to a 
high temperature in an iron retort. The carbon, at the 
high temperature, is able to take the oxygen from the 
potash, forming carbon monoxide, which escapes as a 
gas, while the potassium, which is volatile at a red 
heat, distils over. The preparation of this metal is at- 
tended with many difficulties, and requires special pre- 
cautions, as the vapor of potassium not only takes fire 
when brought in contact with the air, but decomposes 
water, combining with the oxygen and liberating hy- 
drogen, hence the metallic vapor must be cooled by 
rock oil or naphtha, which contains no oxygen. The 
metal thus x>repared must be distilled a second time, in 
order to purify it and free it from a black, explosive 
compound, which invariably forms in the original prep- 
aration. 

Potassium, thus prepared, is a bright, silver-white 
metal, which can be easily cut with a knife at the or- 



POTASSIUM 223 

dinary atmosphere temperature, it is brittle at 0°, and 
melts at 62°. 5, and does not become pasty before melt- 
ing. When heated to a temperature below red heat, 
potassium sublimes, yielding a fine, green-colored va- 
por. This metal rapidly absorbs oxygen when exposed 
to the air, and gradually becomes converted into a 
white oxide. Thrown into water, one atom of potas- 
sium displaces one of hydrogen from the water, form- 
ing potassium hydroxide, or potash. This takes place 
with such force that the heat developed is sufficient to 
ignite the hydrogen thus set free, and the Same be- 
comes tinged with the peculiar purple tint character- 
istic of the potassium compounds, while the water at- 
tains an alkaline reaction from the potash which is 
formed. Potassium also combines directly with chlo- 
rine and sulphur, and many other non-metals, evolving 
heat and light. 

The original source of potassium compounds is the 
felspar of the granitic rocks of which the earth is com- 
posed, as these contain from two to three per cent of 
this metal. Up to the present time, this source has not 
been used for the manufacture of the potassium salts, 
as no cheap and easy mode has yet been available for 
separating the potash from the silicic acid, with which 
it is combined in felspar. Plants, however, are able 
slowly to separate out and assimilate the potash from 
these rocks and soils, so that, by burning the plant and 
extracting the ashes with water, soluble potassium salt 
is obtained. This is the crude potassium carbonate, 
called, when purified by re-crystallization, pearl-ash, 
and it is from this substance that a large number of the 
potassium compounds are obtained. Some of the other 
potassium salts, such as the nitrate and chloride, are 



224 ELEMENTARY CHEMISTRY 

found in large quantities in various localities as de- 
posits on the surface, or in the interior, of the earth. 
Potassium chloride occurs in beds, together with rock 
salt. Another inexhaustible source of potassium com- 
pounds is sea-water. 

Potassium and Chlorine Compounds. 
Potassium Chloride: Formula — ECL 

This salt occurs in certain saline deposits, and also 
exists in large quantities in sea-water. It crystallizes 
in cubes like sodium chloride, and is now much em- 
ployed for the preparation of other potassium salts. 

Potassium and Iodine Compounds. 
Potassium Iodide: Formula — EI. 

A very soluble salt, crystallizing in cubes, obtained 
by dissolving iodine in solution of caustic potash, and 
evaporating and igniting the solid mass to redness, this 
salt is much used as a medicine. 

Potassium and Oxygen Compounds. 

Potassium combines with oxygen in three propor- 
tions, forming three well-defined oxides of the formula, 
potassium monoxide, K 2 0, potassium dioxide, K 2 2 , 
and potassium tetroxide, K 2 4 . 

Potassium monoxide is obtained by allowing thin 
pieces of the metal to oxidize in dry air. It is a grey- 
ish-white, brittle substance, which melts a little above 
red heat, and volatilizes only at a very high tempera- 
ture. This oxide combines with water with evolution 
of great heat, producing potassium hydroxide, or pot- 






POTASSIUM 225 

ash, from which water cannot again be separated by 
heat. 

The dioxide and tetroxide are produced when potas- 
sium is oxidized at high temperatures. 

Potassium, Oxygen and Hydrogen Compounds. 
Potassium Hydroxide or Caustic Potash: Formula — 

KOH. 

Caustic Potash is prepared by boiling one part of 
potassium carbonate with twelve parts of water, and 
adding slaked lime prepared from two-thirds part of 
quicklime. In this reaction calcium carbonate (chalk) 
is formed, which falls to the bottom as a heavy pow- 
der, caustic potash remaining in solution. The clear 
liquid, which should not effervesce on addition of an 
acid, is evaporated in a silver basin to dryness, fused 
by exposure to a stronger heat and cast into sticks in 
a metallic mould. Thus prepared caustic potash is a 
White substance, soluble in half its weight of water, 
and acts as a powerful cautery, destroying the skin. 
It is largely used in the arts and manufactures for 
soap-making, and is employed in the laboratory for 
various purposes. 

Potassium, Carbon and Oxygen Compounds. 
Potassium Carbonate: Formula — K 2 C0 3 . 

This salt receives the commercial name of potash, or 
pearl-ash, and is exported in large quantities from 
Russia and America. The crude substance is prepared 
by boiling out the ashes of plants with water, and 
evaporating the solution to dryness, a pure salt may 
be afterwards obtained by separating the impurities by 



226 ELEMENTARY CHEMISTRY 

crystallization. The leaves and small twigs of plants 
contain more potash than the stems and large branches. 
Potassium carbonate can be obtained perfectly pure by 
heating pure potassium tartrate to redness, and sepa- 
rating the carbonate formed by dissolving in water. 
This salt absorbs water from the air, or is deliquescent, 
and is, therefore, very soluble in water, it also turns 
red litmus blue, or possesses a strongly alkaline reac- 
tion. 

Potassium, Chlorine and Oxygen Compounds. 
Chlorate of Potash: Formula— KCIO.. 

The action of chlorine on potash and the production 
of this salt is done on a large scale by decomposing 
calcium chlorate, made by saturating hot milk of lime 
with excess of chlorine, by means of potassium chlo- 
ride. Potassium chlorate being but slightly soluble in 
cold water separates out in large tubular crystals, while 
the soluble calcium chloride remains dissolved. This 
salt is largely used in medical practice. 



Potassium, Hydrogen, Carbon and Oxygen Compounds. 
Bicarbonate of Potash: Formula — KHC0 3 . 

This substance is formed when a current of carbonic 
acid gas is passed through a strong solution of potas- 
sium carbonate. It may be considered as dibasic car- 
bonic acid, in which one atom of hydrogen is replaced 
by one of potassium. It is a white salt, not so soluble 
as potassium carbonate, the solution is nearly neutral 
to test paper. 



POTASSIUM 227 



Potassium, Nitrogen and Oxygen Compounds. 
Potassium Nitrate or Saltpeter: Formula — KN0 3 . 

This important salt occurs as an efflorescence on the 
soil of several dry tropical countries. It may be arti- 
ficially prepared by the process of nitrification, in 
which animal matter (containing nitrogen) is exposed 
in heaps, mixed together with wood-ashes and lime to 
the action of the air, the organic matter gradually un- 
dergoes oxidation, nitric acid being formed, and this 
unites with the lime and the potash to form nitrates. 
The salt is obtained from both of these sources by boil- 
ing out the soil or deposit with water, adding potas- 
sium carbonate to decompose the nitrate of calcium, 
and allowing the nitre to crystallize out. It dissolves 
in seven parts of water at 15°, and in its own weight 
of hot water. It contains nearly half its weight of oxy- 
gen, with which it parts on heating with carbon or 
other combustible matter. For this reason, nitre is 
largely used in the manufacture of gunpowder and 
fireworks. 

. Potassium, Sulphur and Oxygen Compounds. 
Sulphate of Potash : Formular— K 2 S0 4 . 

Sulphate of Potash is contained in the ashes of both 
sea and land plants, and is only slightly soluble in 
water. A second sulphate termed hydrogen potassium 
sulphate, HKS0 4 , or bisulphate of potash, is a soluble 
salt obtained in the process of the manufacture of ni- 
tric acid. 



228 ELEMENTARY CHEMISTRY 

PHOSPHORUS. 

Symbol— P. Atomic Weight— 30.77. 

Phosphorus does not occur free in nature, but is 
found in combination with oxygen and calcium in large 
quantities in the bodies, and especially the bones, of 
animals, in the seeds of plants, and also as the min- 
erals phosphorite and apatite. When bones are burnt, 
a white solid mass is left behind, this is called calcium 
phosphate (phosphate of lime). Animals obtain the 
phosphate necessary for the formation of their tissues, 
from plants. Plants, again, draw their supply from 
the soil, whilst soils derive their phosphates from small 
quantities existing in the oldest granite rocks, by the 
disintegration of which the fertile soils have been pro- 
duced. Phosphorus appears also to be a very neces- 
sary ingredient in the brain and other centers of the 
nervous action. 

Phosphorus is prepared from powdered bone-ash, by 
mixing it with two-thirds of its weight of sulphuric 
acid and 15 to 20 parts of water. The sulphuric acid 
decomposes the bone-ash, forming calcium sulphate, or 
gypsum, which separates out as a white insoluble pow- 
der, while the greater part of the phosphorus in the 
bones comes into solution in combination with calcium, 
oxygen, and hydrogen, forming calcium hydrogen phos- 
phate, a salt commonly known as superphosphate of 
lime. The liquid is drawn off clear, evaporated down 
to a syrup, and then mixed with powdered charcoal, 
dried, and heated to redness in an earthenware retort, 
the neck of which dips under water. Half the phos- 
phorus is liberated together with carbon monoxide, and 



PHOSPHORUS 229 

distils over, collecting under the water in yellow drops, 
while the other half remains behind in the retort as 
calcium pyrophosphate. 

In order to purify the phosphorus thus prepared, it 
may again be distilled, or pressed when melted under 
hot water through leather, it is then cast into sticks 
and kept under cold water. Phosphorus is an exceed- 
ingly inflammable and oxidizable substance, and re- 
quires great care in its preparation. It is manufac- 
tured on a very large scale for making the composition 
for the tips of matches. Phosphorus is a slightly yel- 
low semi-transparent solid, resembling white wax both 
in appearance and consistency, but at low temperature 
it becomes brittle. It melts at 44°, forming a transpar- 
ent liquid. It boils at 290°, giving rise to a colorless 
gas. In the air it gives off white fumes, emitting a 
pale phosphorescent light in the dark, whence its name, 
it is then undergoing a slow combustion, the w T hite 
fumes consisting of phosphorus trioxide. At a tem- 
perature very little above its fusing point, phosphorus 
takes fire in the air, entering into active combustion, 
and forming phosphorus pentoxide,. or phosphoric an- 
hydride. The ignition of phosphorus takes place by 
slight friction, or by a blow, and even the heat of the 
hand may cause this substance to ignite, hence great 
care must be taken in handling phosphorus, and it 
should always be cut under water. Phosphorus does 
not dissolve in water, alcohol, or ether, but it is slight- 
ly soluble in oils, and very readily soluble in carbon 
disulphide. 

If yellow phosphorus be exposed to a temperature 
of about 240° for some hours in an atmosphere in- 
capable of acting chemically on it, such as hydrogen 



230 ELEMENTARY CHEMISTRY 

or carbon dioxide, it is found to have undergone a 
very remarkable change, being wholly converted into 
a dark red opaque substance, altogether insoluble in 
carbon disulphide. The weight of red substance pro- 
duced is exactly equal to that of yellow phosphorus 
used. This is called red, or amorphous phosphorus, and 
differs much in its properties from the yellow modifi- 
cation, especially in its inflammability, as it does not 
take fire in the air until heated to above 260°, when it 
becomes reconverted into the ordinary form, and burns 
with the formation of phosphorus pentoxide. The sud- 
den conversion of yellow into red phosphorus can be 
shown by heating a small piece of ordinary phosphorus 
in a dry tube with a mere trace of iodine, combination 
at once occurs, a small trace of volatile phosphorus 
iodide is formed, and the remainder of the phosphorus 
is converted into the red modification. The red or 
amorphous modification of phosphorus can also be ob- 
tained in a crystallized form by heating red phosphorus 
in a tube with metallic lead. The phosphorus dissolves 
in the melted lead, and on cooling separates out in crys- 
tals, which possess a bright black metallic lustre. 

Phosphorus appears luminous in the dark, when in 
contact with moist air, and it evolves fumes possessing 
a strong garlic-like smell. These fumes are poisonous, 
producing phosphorus-necrosis, a disease in which the 
bones of the jaw are destroyed, and one by which scrof- 
ulous subjects are the most easily affected. The lumi- 
nosity of phosphorus in the air depends upon its slow 
oxidation, with formation of phosphorous acid. In this 
act of combination so much heat is evolved, that if a 
large piece of phosphorus be allowed to lie exposed 
to the air it at last melts and then takes fire. The lumi- 



PHOSPHORUS 231 

nosity and oxidation of phosphorus are best seen by 
pouring a few drops of the solution of this body in 
carbon bisulphide on to a piece of filter paper and al- 
lowing the solution to evaporate. In the dark the pa- 
per soon begins to exhibit a bright phosphorescence, 
and after a short time the phosphorus takes fire and 
burns. It was formerly believed that phosphorus be- 
comes luminous in gases upon which it can exert no 
chemical action, such as hydrogen or nitrogen. This is 
not so, the luminosity which has been observed in these 
cases being due to the presence of traces of oxygen. 
From these facts it would naturally be inferred that 
phosphorus must be more luminous in pure oxygen 
than in air, this is not the case. At temperatures be- 
low 20° phosphorus is not luminous in pure moist oxy- 
gen, indeed it may be preserved for many weeks in this 
gas without undergoing the slightest oxidation. If, 
however, the gas be diluted by admixture with another 
indifferent gas, or if it be rarified, the phosphorescence 
is at once observed. The phenomenon can be very beau- 
tifully shown by placing a stick of phosphorus in a long 
tube a, Fig. 36, closed at one end and open at the 
other, and partly filled with mercury, into which some 
pure oxygen is brought. The open end of the tube is 
connected by a rubber tube with the vessel b containing 
mercury, so that, by raising or lowering the vessel, the 
pressure on the gas can be regulated. If the pressure 
be so arranged that it does not amount to more than 
one-fifth of an atmosphere, the phosphorus will be seen 
to be brightly luminous in the dark. If the pressure be 
then gradually increased, the light will become less 
and less distinct, until, when the level of the mercury 
is the same in both vessels, the luminosity has entirely 



232 



ELEMENTARY CHEMISTRY 



ceased. The phosphorescence can, however, at once be 
brought back again by lessening the pressure. 




Fig. 36. 



SILICON 233 

SILICON. 

Symbol— Si. Atomic Weight— 28.2. 

Silicon, next to oxygen, is the most abundant ele- 
ment known. It does not occur, however, in the free 

state, but always combined with oxygen to form silicon 
dioxide, or silica. Silicon dioxide exists nearly pure 
in quartz or rock crystal, in flint, sand, and in a variety 
of minerals. Silicon also occurs combined with metals 
and oxygen, forming metallic silicates, and of these 
the greater part of almost all known rocks, especially 
the primary rocks, is composed. 

In order to obtain silicon in the free state, a com- 
pound of this substance with fluorine and potassium, 
potassium silico-fluoride, is heated with metallic potas- 
sium, a violent reaction occurs, and when the contents 
of the tube in which the decomposition was effected are 
put into water, silicon is left undissolved in the form 
of a brown amorphous powder. Silicoa can be ob- 
tained in three different modifications : Amorphous, 
graphitoidal, and crystalline. The graphite form of 
silicon is prepared by heating the brown amorphous 
powder to a high temperature, when the mass con- 
tracts, and becomes much more dense. Crystalline 
silicon is best obtained by fusing the mixture which 
gives brown silicon with zinc. On cooling the mass, 
crystals of silicon are found to be deposited on the 
zinc, which latter can easily be removed by solution 
in an acid. Silicon thus obtained is hard enough to 
scratch glass, it may be fused at a temperature between 
the melting point of cast-iron and steel. 

Pure silicon may best be prepared by heating a mix- 



234 



ELEMENTARY CHEMISTRY 



ture of powdered quartz and magnesium, in the calcu- 
lated quantities with a quarter of its weight of cal- 
cined magnesia in a fire-clay crucible, first at 300° to 
400° in order thoroughly to dry the material, and then 
for a few minutes at a red heat, when a vigorous reac- 
tion takes place. 

After the action has ceased, the cooled mass is treated 
with hydrochloric acid to dissolve out the magnesia, 
then warmed several times alternately with hydro- 
fluoric and sulphuric acids to remove unchanged silica, 
finally washed with water and dried by heating in a 
stream of dry hydrogen. By employing precipitated 
silica and pure magnesium a very pure form of silicon 
may be obtained. An impure form, which may be ad- 
vantageously used for the preparation of many of its 
derivatives, can be rapidly obtained by heating 40 
grams of dry powdered white sand with 10 grams of 
magnesium powder in a wide test-tube. The reaction 
is a tolerably vigorous one, and if carried out Avith 
precipitated silica is accompanied by a brilliant flash 
of light. 

Pure silicon prepared as above is an amorphous ma- 
roon-colored powder. When heated in air it oxidizes 
superficially, and it takes fire in oxygen at about 400°, 
burning brilliantly to silica, whilst it decomposes steam 
slowly at a dull red heat forming silica and hydrogen. 

Amorphous silicon dissolves in many molten metals, 
uniting with the metal in some cases, as with iron, 
nickel, magnesium, and copper, to form metallic sil- 
icides, whilst in other cases, as, for example, with 
aluminium, zinc, and silver, no combination takes place, 
but on cooling the mass, silicon separates out in the 
crystalline form. 



SILICON 235 

A different variety of amorphous silicon, which pos- 
sesses remarkable reducing properties, is obtained 
when sparks are passed through liquid silico-ethane, 
the silicon, prepared in this way, reduces neutral solu- 
tions of potassium permanganate in the cold, and cop- 
per sulphate, gold and mercuric chlorides on boiling. 

Silicon may also be obtained crystalline either in six- 
sided plates or long needles, and these forms are some- 
times distinguished as graphitoidal and adamantine 
silicon. The crystalline modification of silicon is pre- 
pared by heating metallic aluminium to redness with 
three times its weight of potassium silico-fluoride in a 
fire-clay crucible, for about half-an-hour, when the fol- 
lowing reaction takes place : 

3K 2 SiF e +4Al=6KF+4AlF 3 +3Si. 
The temperature is then raised to about 1,000° and 
the silicon dissolves in the excess of aluminium, and 
on cooling crystallizes out in thin plates, w^hich are 
separated from the aluminium and potassium fluorides 
by washing with hydrochloric and hydrofluoric acids. 

Silicon and Carbon Compounds. 
Silicon Carbide: Formula— SiC. 

Silicon carbide is formed when a mixture of coke, 
sand and salt is fused in an electric furnace employing 
carbon terminals, and may be obtained pure by fusing 
silicon with the requisite amount of carbon in the elec- 
tric furnace. It is also formed by the action of car- 
bon on iron or calcium silicide, by fusing calcium car- 
bide with silica in the electric furnace, and by the 
union of the vapors of carbon and silicon in the electric 
furnace, needles of the carbide being deposited. The 



236 ELEMENTARY CHEMISTRY 

crystals are colorless or sapphire blue and are usually 
found as six-sided plates, but when formed by the com- 
bination of the two vapors they are prismatic. Silicon 
carbide is not oxidized by oxygen at 1,000° and is not 
attacked by sulphur vapor, by fused potassium nitrate, 
or by any acid. It is completely decomposed by chlo- 
rine at 1,200°, and on fusion with lead chromate it is 
gradually oxidized, while fused caustic potash slowly 
converts it into potassium carbonate and silicate. 

The crude material, obtained by the method first de- 
scribed, is known as carborundum and is used as a cut- 
ting and polishing agent. It is manufactured in large 
quantities at Niagara Falls. 

Silicon and Oxygen Compounds. 

Silicon Dioxide or Silica: Formula — Si0 2 . Molecular 

Weight^-59.96. 

Silicon forms only this oxide, which is an extremely 
important constituent of our planet. It is found not 
only in the mineral but also in the vegetable and ani- 
mal kingdoms, existing in large quantity in the glassy 
straw of the cereals and of bamboos, and in the feathers 
of certain birds, which have been found to contain as 
much as 40 per cent of this substance, while silicic 
acid has been found in all forms of connective tissue. 
Vast deposits of pure silica in a very fine state of di- 
vision occur in various parts of Germany, especially 
in Hanover and near Berlin. Large quantities of this 
substance are now used for a variety of purposes, es- 
pecially for the preparation of dynamite, for filtering, 
and as a non-conducting medium for packing steam- 
pipes. In the mineral kingdom it is found in three 
distinct forms. 



SILICON 237 

Silica melts in the oxy -hydro gen flame to a colorless 
glass which may be drawn out in threads. It slowly 
volatilizes when maintained at a high temperature in a 
furnace for some time, and can readily be made to 
boil when heated in an electric furnace. 

When an alkaline solution of silica is heated in a 
sealed tube the glass is attacked and an acid silicate is 
formed from which silica separates out on cooling. If 
the temperature at which the deposition occurs be 
above 180°, the silica separates out as quartz, if below 
this point, it crystallizes out as tridymite, whilst at the 
ordinary temperature of the air it separates in the form 
of a hydrated amorphous mass. 

The various forms of silica are employed in a variety 
of technical processes, especially its application to the 
manufacture of glass and porcelain. The colored varie- 
ties of silica are also largely used as gems and for other 
ornamental purposes. It is also possible to color the 
natural agates artificially. Thus brown or yellow 
agates or chalcedonies when strongly heated are 
changed into ruby carnelians, the yellow oxide of iron 
being thereby changed into the red anhydrous oxide. 
Many agates and chalcedonies are permeable to liquids, 
and in this way they may be artificially colored. This 
fact was known to the ancients and made use of in 
darkening the color of agates. 

Silicon, Oxygen and Hydrogen Compounds. 
Silicic Acid: Formula— Si (OH) 4 . 

If a solution of an alkali silicate, termed soluble 
glass, be acidified with hydrochloric acid, a portion of 
the silicic acid separates out as a gelatinous mass, 
whilst another portion remains in solution. If, on the 



238 ELEMENTARY CHEMISTRY 

other hand, the solution is sufficiently dilute, no pre- 
cipitate will occur, all the silicic acid remaining dis- 
solved. This liquid contains silicic acid, hydrochloric 
acid, and common salt in solution. In order to separate 
the last two compounds from the first, the liquid is 
brought into a flat drum, the bottom of which consists 
of parchment paper, and this dialyser containing the 
liquid is allowed to swim on the surface of a large vol- 
ume of water. The sodium chloride and the excess of 
hydrochloric acid pass through the membrane, whilst 
a clear aqueous solution of silicic acid remains behind. 

This mode of separation was termed dialysis by its 
discoverer. Substances which crystallize, hence termed 
crystalloids, have as a rule the power of passing in 
solution through a membrane such as parchment or an 
animal membrane, whilst substances, such as gum and 
glue, which form jellies and are termed colloids, are 
unable to pass through such a diaphragm. 

A fluid mixture of a colloid and a liquid is termed a 
sol, a watery mixture being called a hydrosol, whilst a 
firm mixture of a colloid and a liquid is named a "gel," 
thus silicic acid jelly is a hydrogel of silicic acid 

These so-called colloidal solutions, which are thought 
by many to be suspensions and not true solutions, have 
the power of polarizing transmitted light. 

By dialysis an aqueous solution of pure silicic acid 
may be obtained which contains 5 per cent of silica, 
and this may be concentrated by boiling it in a flask 
until it reaches a strength of 14 per cent. When heated 
in an open vessel, such as an evaporating basin, it is 
apt to gelatinize round the edge, after which the whole 
solidifies. The solution of silicic acid thus prepared 
has a feebly acid reaction and is colorless, limpid, and 



SILVER 239 

tasteless. On standing for a few days the solution 
gelatinizes to a transparent jelly. This coagulation 
is retarded by the presence of a few drops of hydro- 
chloric acid, or of caustic alkali, but it is brought 
about by even the smallest traces of an alkali car- 
bonate. 

If the clear solution be allowed to evaporate in a 
vacuum at 15° a transparent glass-like mass remains 
behind, which, when dried over sulphuric acid, pos- 
sesses approximately the formula. H o Si0 3 =Si0 o +II o 0. 
This has been termed metasilicic acid, and an acid of 
the same composition has been obtained by dehydrating 
precipitated gelatinous silicic acid with 90 per cent 
alcohol. 

SILVER. 

Symbol— Ag. Atomic Weight— 107.13. 

Silver is found in the native state, as well as com- 
bined with sulphur, antimony, chlorine, and bromine. 
It is also contained in small quantities in galena, and 
it can be extracted with profit from the lead prepared 
from this ore, even when the lead contains only two 
or three ounces of silver to the ton. The method adopt- 
ed for the extraction of the silver depends upon the 
fact that the whole of the silver can be concentrated 
into a small portion of lead, by crystallization, metallic 
lead free from silver separates out in crystals, and a 
rich alloy is left. When this reaches the concentration 
of 300 ounces of silver to the ton, the alloy undergoes 
the operation of cupellation, in which the mixture is 
melted in a furnace on a porous bed of bone-earth, and 
a blast of air blown over the surface. The lead oxi- 
dizes, and the oxide (litharge) fuses, and partly runs 



240 ELEMENTARY CHEMISTRY 

away and partly sinks into the porous bed of the fur- 
nace, whilst the silver remains behind in the metallic 
state. 

For the extraction of silver from the other ores, a 
process termed amalgamation is employed, in which 
mercury is used to dissolve the metallic silver. The 
argentiferous ores in which the silver occurs in com- 
bination with sulphur, are worked in a different man- 
ner. The ore is roasted in a furnace with common salt, 
by which means the silver sulphide is converted into 
chloride, the mixture is then placed in casks made to 
revolve, and scrap-iron and water are added. The 
iron reduces the silver to the metallic state, and when 
this is fully accomplished, metallic mercury is added. 
This forms a liquid amalgam with the silver, and gold, 
if any be«present, and by distilling the mercury off, the 
silver is obtained in an impure state. Silver possesses 
a bright white color and a brilliant lustre, which it does 
not lose in pure air at any temperature, but when melt- 
ed in the air it possesses the singular power of absorb- 
ing mechanically a large volume (twenty-two times its 
bulk) of oxygen, this gas it again gives out on solidify- 
ing, a phenomenon known as the spitting of silver. 

Silver is probably the best conductor of heat and 
electricity known, and is extremely ductile, one 
gramme of metal can be drawn out into a wire of 
2,600 metres in length. Sulphur combines at once with 
silver, forming a black sulphide, silver articles long 
exposed to the air tarnish from this cause. Silver is 
easily soluble in nitric acid, the nitrate being formed 
and nitric oxide gas being evolved. 



SILVER 241 

Silver and Chlorine Compounds. 
Silver Chloride: Formula — AgCl. 

Of the insoluble salts, the Silver chloride is the most 
important. This salt occurs in nature, and is then 
known as horn silver, and is precipitated as a white 
curdy mass when solutions of a chloride and a silver 
salt are brought together. "When exposed to sun or 
day-light, the white chloride becomes tinted a purple 
color, w^hieh increases in shade as the action of light 
continues. This coloration arises from a partial de- 
composition of the salt, a small quantity of argentous 
chloride and free hydrochloric acid being formed. 

Silver and Oxygen Compounds. 

Silver suboxide, Ag 4 0, Silver monoxide, Ag 2 and Sil- 
ver dioxide, Ag 2 2 . 

Silver suboxide : Formula — Ag 4 0. 

Silver forms three compounds with oxygen. The 
first of these is called Silver suboxide, it is a black 
powder which readily undergoes decomposition. 

Silver monoxide: Formula — Ag 2 0. 

The second, termed Silver monoxide, is obtained in 
the form of a brown precipitate, when caustic potash 
is added to a solution of silver nitrate. From this 
oxide, which is decomposed into metal and oxygen on 
heating, the ordinary silver salts may be derived by 
dissolving it in acids. 



242 ELEMENTARY CHEMISTRY 

Silver dioxide: Formula — Ag 2 2 . 

The third oxide is called Silver dioxide, and is ob- 
tained as a black powder by the action of ozone on 
metallic silver. 

Silver, Nitrogen and Oxygen Compounds. 
Nitrate of Silver: Formula — AgN0 2 . 

Silver nitrate is the most important soluble salt of 
silver. It is obtained in the form of large transparent 
tubular crystals on evaporating a solution of silver in 
nitric acid, and is soluble in its own weight of cold 
and half its weight of hot water, and in four parts of 
alcohol. Silver nitrate fuses easily on heating, and 
when cast into sticks goes by the name of lunar caus- 
tic. This salt undergoes decomposition when exposed 
to the sunlight in contact with organic matter, and a 
black substance, consisting of the suboxide, is formed. 
It is employed in the manufacture of an indelible ink 
for marking linen and other fabrics, and for photo- 
graphic purposes. 

SODIUM. 

Symbol— Na. Atomic Weight— 22.87. 

This metal was discovered by the decomposition of 
soda with the galvanic current. It can be procured 
more easily than potassium by reducing the carbonate 
in presence of carbon, and is now manufactured in 
large quantities for the preparation of other metals, 
especially magnesium and aluminium. The apparatus 
employed for the preparation of this metal is the same 
as that used for potassium. The metal distills over 



SODIUM 243 

when condensed, and drops into rock oil. Sodium is 
a silver-white metal, soft at ordinary temperatures, 
and melting at 95.6°, it volatilizes below a red heat, 
yielding a colorless vapor. When thrown upon water 
it floats, and rapidly decomposes the water with dis- 
engagement of hydrogen, soda being formed. If the 
water be hot or be thickened w T ith starch, the globule 
of metal becomes so much heated as to enable the hy- 
drogen to take fire. The compounds of sodium are 
very widely diffused, being contained in every parti- 
cle of dust. They exist in enormous quantities in the 
primitive granitic rocks, but they are most readily ob- 
tained from sea-water, which contains nearly three per 
cent of sodium chloride (common or sea salt). Sodium 
carbonate was formerly obtained from the ashes of 
sea-plants or kelp, as potassium carbonate is still pre- 
pared from the ashes of land plants, but at present the 
sodium carbonate is altogether manufactured, on a 
large scale, from sea-salt. 

Sodium, Carbon and Oxygen Compounds. 
Sodium Carbonate or Soda-ash: Formula — Na 2 C0 3 . 

This substance, known in commerce as soda-ash, is 
manufactured on an enormous scale, and used for glass- 
making, soap-making, bleaching, and various other 
purposes in the arts. Formerly it was prepared from 
barilla or the ashes of sea-plants, but now it is wholly 
obtained from sea-salt by a series of chemical decom- 
positions and processes. 



244 ELEMENTARY CHEMISTRY 

Sodium and Chlorine Compounds. 
Chloride of Sodium or Common Salt. Formula — NaCl. 

It is from this salt that almost all the other sodium 
compounds are prepared. Sodium chloride occurs in 
thick beds in various, parts of the world. It is like- 
wise prepared from sea-water by evaporation or by 
freezing, and from certain brine springs by evapora- 
tion. When slowly deposited sodium chloride crystal- 
lizes in regular cubes. It is soluble in about two and 
a half parts of water at 15 °, and does not dissolve sen- 
sibly more in hot than in cold water. 

Sodium and Oxygen Compounds. 
Sodium Oxide, Na 2 and Sodium Dioxide, Na 2 2 . 

Sodium Oxide: Formula — Na 2 0. 

Sodium Oxide is formed when sodium is oxidized 
in dry air or oxygen at a low temperature, a white 
powder being formed. This takes up moisture with 
great avidity, # forming sodium hydroxide, or soda, 
from which water cannot again be separated by heat 
alone, but which can again be converted into the oxide 
by heating with sodium. 

Sodium Dioxide: Formula — Na 2 2 . 

Sodium Dioxide is a yellowish-white powder, which 
is formed when sodium is heated in oxygen to 200° 
C. It is soluble in water, but the solution readily de- 
composes, giving off oxygen and leaving sodium hy- 
droxide. 



SODIUM 245 

Sodium, Hydrogen and Oxygen Compounds. 
Sodium Hydroxide or Caustic Soda: Formula — NaHO. 

Sodium Hydroxide is a white solid substance, fusi- 
ble below a red heat, and less volatile than the corre- 
sponding potassium compound. It is very soluble in 
water, acts as a caustic, is powerfully alkaline and is 
largely used in soap-making. The manufacture of 
solid caustic soda is carried on on a large scale, by 
boiling lime and sodium carbonate together with water, 
and evaporating down the clear solution. 

Sodium, Hydrogen, Carbon and Oxygen Compounds, 
Sodium Bicarbonate: Formula — NaHC0 8 . 

Bicarbonate of Soda is obtained by exposing the 
crystallized carbonate in an atmosphere of carbonic 
acid gas. It is a white crystalline powder, which on 
heating is readily converted into sodium carbonate. The 
bicarbonate is chiefly used in medicine, and for the 
production of effervescing drinks. 

If sodium carbonate be added to a solution of tri- 
hydrogen phosphate, effervescence will at once ensue 
from the liberation of carbonic acid, and if the car- 
bonate be added until the solution ceases to redden 
litmus paper, a salt will be obtained on evaporation 
which crystallizes in large transparent prisms. This 
is common sodium phosphate, its composition is repre- 
sented by the symbol Na 2 HP0 4 , with twelve molecules 
of water of crystallization. If caustic soda be added 
to a solution of this common phosphate, a salt termed 
the subphosphate, crystallizes out in small needles on 
evaporation, the composition of this salt is Na 3 P0 4 , 



246 ELEMENTARY CHEMISTRY 

with twelve atoms of water of crystallization. And if 
phosphoric acid be added to a solution of common 
phosphate, the so-called sodium superphosphate is 
formed, NaH 2 P0 4 . 

Sodium, Nitrogen and Oxygen Compounds. 
Sodium Nitrate or Saltpeter: Formula— NaN0 2 . 

Sodium nitrate is found in large beds in Peru and 
Chili, and is termed Chili saltpeter. It is imported in 
large quantities and used as a manure, and also in the 
preparation of nitric acid, being cheaper than nitre. 
For this latter purpose a hot concentrated solution of 
this salt is mixed with a hot saturated solution of 
potassium chloride, on cooling, potassium nitrate sep- 
arates out in crystals, and sodium chloride remains in 
solution. 

Sodium, Sulphur and Oxygen Compounds. 

Sulphate of Soda or Glauber's Salts: Formula— 

Na 2 SO 4 +10H 2 O. 

Sodium sulphate is known in commerce as Glauber's 
salts, and in the anhydrous state as salt-cake. It oc- 
curs in the water of many mineral springs, and is used 
in medicine, while as salt-cake it is employed in large 
quantities in glass manufacturing. 

SULPHUR. 
Symbol— S. Atomic Weight—31.8. 

Sulphur occurs both free and combined in nature. It 
is found free in certain volcanic countries, especially 
in Sicily, Iceland and Mexico, and occurs crystallized 
in yellow transparent crystals. It exists in combina- 



SULPHUR 247 

tion with many metals, forming compounds termed 
sulphides, which constitute the common ores from 
which the metals are usually obtained, thus lead sul- 
phide, or Galena, zinc sulphide, or Blende and copper 
sulphide, are the substances from which those metals 
are generally procured. Sulphur also is found in na- 
ture combined with metals and oxygen, to form a class 
of salts called sulphates. Of these, calcium sulphate 
or gypsum, barium sulphate or heavy spar, sodium 
sulphate or Glauber's salt, occur in the largest quan- 
tity, Sulphur likewise occurs combined with hydro- 
gen, as a gas called Sulphuretted Hydrogen, H 2 S, in 
the waters of certain springs. In order to obtain pure 
sulphur, the mineral containing the crude substance 
mixed with earthy impurities is heated in earthenware 
pots, the sulphur distills over in the form of vapor, 
which is condensed in similar pots placed outside the 
furnace. AYhen brought to this country, the sulphur 
thus obtained is refined or purified by subjecting it to 
a second distillation. If the vapor of sulphur is quick- 
ly cooled below its melting point, it solidifies in the 
form of a fine crystalline powder called Flowers of 
Sulphur, exactly as aqueous vapor, when cooled down 
below the freezing point of water, deposits as snow. 
When sulphur is gently heated, it melts, and may be 
cast into sticks, and is then known as brimstone or 
roll sulphur. 

Sulphur exists in three modifications. The first is 
that in which sulphur crystallizes in nature, and the 
other two are obtained by melting sulphur. If melted 
sulphur be allowed to cool slowly, it crystallizes in 
long, transparent, needle-shaped, prismatic crystals, 
which are quite different in form from the natural 



248 ELEMENTARY CHEMISTRY 

crystals of sulphur. These transparent crystals become 
opaque after exposure to the air for a few days, owing 
to each crystal splitting up into several crystals of 
the natural or permanent form. The third allotropic 
modification of sulphur is obtained by pouring melted 
sulphur heated to 230° into cold water. The sulphur 
thus forms a soft tenacious mass resembling rubber. 
This form of sulphur is, however, not permanent, in a 
few hours, at the temperature of the air, the mass as- 
sumes the ordinary brittle form of the element, while, 
if heated to 100°, it instantly changes to the brittle 
form, and thereby evolves so much heat as to raise its 
temperature up to 111°. These peculiar modifications 
become apparent when sulphur is heated. Sulphur 
melts, to begin with, at 115°, forms an amber-colored 
mobile liquid, as the temperature rises, the liquid be- 
comes dark-colored, and attains the consistency of thick 
molasses, so that at about 230° it can scarcely be 
poured out of the vessel, heated above 250°, it again 
becomes fluid, and remains as a dark reddish-black 
colored thin liquid, until the temperature rises to 490°, 
when it begins to boil, and gives off a red-colored vapor. 
Sulphur is an inflammable substance, and when heat- 
ed in the air or in oxygen burns with a bluish flame, 
combining with the oxygen to form sulphur dioxide, 
often, called sulphurous acid, which is given off as a 
gas, possessing the peculiar and well-known suffocating 
smell which is evolved when a common match is burnt. 
Sulphur combines directly with chlorine, carbon, and 
most other elements, while many metals burn in sul- 
phur vapor as they do in oxygen gas, uniting to form 
sulphides. Sulphur is insoluble in water and most 
organic liquids, but both the natural variety and the 



SULPHUR 



249 



other crystalline or prismatic variety dissolve freely 
in carbon disulphide, while the tenacious form of sul- 
phur is insoluble in this liquid. When deposited from 
solution in carbon disulphide, sulphur crystallizes in 
the ordinary natural form. 




Fig. 37. 



Commercial sulphur contains about 3 per cent of 
earthy impurities which can be removed by distilla- 
tion, an arrangement for this purpose being shown in 
Fig. 37. The sulphur is melted in an iron pot M and 
runs from this by means of a tube into the iron retort 
G where it is heated to the boiling point, the vapor of 



250 ELEMENTARY CHEMISTRY 

the sulphur then passes into the large chamber A which 
has a capacity of 200 cubic metres. In this chamber 
the sulphur is condensed, to begin with, in the form 
of a light yellow powder termed flowers of sulphur, 
just as aqueous vapor falls as snow when the tem- 
perature suddenly sinks below 0°. After a time the 
chamber becomes heated above the melting point of 
sulphur, and then it collects as a liquid which can be 
drawn off by means of the opening 0. It is then cast 
in slightly conical wooden moulds, and is known as 
roll sulphur or brimstone. It is frequently also al- 
lowed to cool in the chamber and then obtained in 
large crystalline masses, known to the trade as block 
sulphur. 

Another and much more important source from which 
sulphur is now obtained is the residue or waste in the 
soda manufacture, this consists of calcium sulphide 
mixed with chalk, lime, and alkali sulphides. The sul- 
phur which this material contains was formerly alto- 
gether wasted, now it is economically regained in the 
alkali works on a very large scale. For this purpose 
the waste, which consists mainly of calcium sulphide, 
CaS, is treated in the presence of water with carbonic 
acid gas, calcium carbonate being thus produced whilst 
sulphuretted hydrogen is evolved. 

The gas thus obtained is then mixed with a quantity 
of air sufficient to supply the amount of oxygen re- 
quired and the mixture passed through a kiln, where 
it comes in contact with heated ferric oxide and the 
above reaction takes place, the sulphur formed being 
condensed in cooling chambers, and recovered in the 
pure state. A less pure sulphur is obtained in the 
same way from the sulphuretted hydrogen given off 



SULPHUR 251 

during the distillation of the ammoniacal liquor from 
gas works. 

The simplest mode of detecting sulphur in a com- 
pound is to mix the substance with pure sodium car- 
bonate, and fuse it before the blowpipe on charcoal, or, 
to avoid the introduction of sulphur from the gas 
flame, to mix the substance with sodium carbonate and 
charcoal, and heat in a small closed crucible. Sodium 
sulphide is thus formed, and this may then be recog- 
nized by bringing the fused mass on to a silver coin and 
adding water. The smallest quantity of sulphur can 
thus be recognized by the formation of a brown stain 
of silver sulphide. Sulphur is almost always quanti- 
tatively determined as barium sulphate. If the body is 
a sulphide, as for instance, pyrites, it is finely pow- 
dered, and either fused with a mixture of sodium car- 
bonate and nitre, the fused mass dissolved in water, 
and the filtrate, after acidifying by hydrochloric acid, 
precipitated with barium chloride, or the sulphide is 
oxidized with a mixture of nitric and hydrochloric 
acids or fuming nitric acid, the excess of acid re- 
moved by evaporation and barium chloride added, 
whereby insoluble barium sulphate is formed, and this, 
after washing and drying, is ignited and weighed. A 
further method is to fuse the mineral with 2 parts of 
caustic soda and 4 parts of sodium peroxide, the melt 
being then acidified and precipitated with barium 
chloride as above. 

Sulphur and Oxygen Compounds. 

Sulphur Dioxide, S0 o and Sulphur Trioxide, S0 3 . 



252 ELEMENTARY CHEMISTRY 

Sulphur Dioxide: Formula— S0 2 . Molecular Weight 
—63.59. Density— 31.8. 

Sulphur dioxide is formed not only by the combus- 
tion of sulphur, but also by the action of certain metals, 
such as copper, mercury, or silver, on concentrated 
sulphuric acid. 

Sulphur dioxide is easily prepared for laboratory 
use. For this purpose a flask is half filled with copper 
turnings or fine copper foil, and so much strong sul- 
phuric acid poured in that the copper is not quite cov- 
ered. The mixture is next heated until the evolution 
of gas commences, the lamp must then be removed, as 
otherwise the reaction may easily become too violent, 
and the liquid froth over. 

Pure sulphur dioxide is also produced when sulphur 
and sulphuric acid are heated together, thus 

S+2H 2 S0 4 =3S0 2 +2H 2 0. 

It is also formed by the decomposition of a sulphite, 
such as commercial sodium sulphite, which, when 
treated with warm dilute sulphuric acid, easily evolves 
the gas, thus 

Na 2 S0 3 +H 2 S0 4 =Na 2 S0 4 +H 2 0+S0 2 . 

A convenient method of employing this decomposi- 
tion is to allow concentrated sulphuric acid to drop 
into a saturated solution of sodium bisulphite. 

Sulphur dioxide is made on a large scale for the 
preparation of the sulphites, especially of sodium sul- 
phite and calcium sulphite, which are obtained by 
passing the gas either into a solution of caustic soda 
or into milk of lime. For this purpose charcoal is 
heated together with sulphuric acid, when carbon di- 



SULPHUR 



253 



oxide is evolved, together with sulphur dioxide, but 
the presence of the former compound for tht purpose 
above mentioned is not detrimental, thus 

C-f2H 2 S0 4 =2H 2 0+C0 2 +2S0 2 . 

The sulphur dioxide evolved in the roasting of cer- 
tain metallic ores, which was formerly allowed to pass 
off into the atmosphere, is now frequently utilized for 
the preparation of the sulphites. 



Fig. 38. 

Sulphur dioxide is used in enormous quantities for 
the manufacture of sulphuric acid. For this purpose 
it is chiefly obtained by roasting pyrites. When, how- 
ever, especially pure sulphuric acid is needed, the 
dioxide is prepared by burning pure sulphur. 

Sulphur dioxide is a colorless gas, which occurs in 



254 ELEMENTARY CHEMISTRY 

nature in certain volcanic emanations, as well as in 
solution in volcanic springs. It possesses the well- 
known suffocating smell of burning sulphur. It can 
be collected by downward displacement, like chlorine. 
In order to prepare liquid sulphur dioxide in larger 
quantity, the apparatus Fig. 38 is used. The gas 
evolved by the action of sulphuric acid on copper is 
purified by passing through the wash-bottle, and after- 
wards passes through the spiral glass tube, surrounded 
by a freezing mixture of ice and salt. The liquid 
which condenses and falls into the flask placed be- 
neath may be preserved by sealing the flask hermeti- 
cally where the neck has been drawn out. 

Sulphur Trioxide : Formula— SO r Molecular Weight — 
79.47. Density— 39.73. 

Sulphur trioxide is obtained by passing a perfectly 
dry mixture of sulphur dioxide and oxygen through 
a tube containing heated platinum sponge or platinized 
asbestos as in Fig. 39. The sulphur dioxide is evolved 
in the flask a and is mixed in the wash-bottle, which 
contains strong sulphuric acid, with the oxygen from 
a gas-holder coming in through the tube b. The mix- 
ture next passes through the cylinder e containing 
pumice-stone soaked in strong sulphuric acid in order 
to remove every trace of moisture, and then passes at 
c over the platinized asbestos. As long as this is not 
heated no change is observed, so soon, however, as it 
is gently ignited dense white fumes of the trioxide are 
formed which condense in a receiver d, cooled by a 
freezing mixture, in the form of long white needles. 
In order to obtain these crystals, every portion of the 



SULPHUR 



255 



apparatus must be absolutely dry, if even a trace of 
moisture be present the needles disappear at once, 
liquid sulphuric acid being formed. Instead of plati- 
num, certain metallic oxides, such as copper oxide, 
ferric oxide, and chromic oxide, may be used. 

A much more convenient method of preparing sul- 
phur trioxide for experimental purposes is by the dis- 




Fig. 39. 



tillation of fuming sulphuric acid, which consists of a 
solution of the trioxide in sulphuric acid, and is manu- 
factured in large quantity. 

The trioxide absorbs moisture rapidly from the at- 
mosphere, and evolves dense white fumes on exposure 
to the air, and when brought in contact with anhy- 
drous baryta, combines with it to form barium sul- 
phate, BaS0 4 , with such energy that the mass becomes 
incandescent. 



256 ELEMENTARY CHEMISTRY 

Sulphur trioxide is now manufactured in very large 
quantities by the combination of sulphur dioxide and 
oxygen in presence of catalytic agents, such as plati- 
num or ferric oxide. 

\ 

TIN. 

Symbol— Sn. Atomic Weight— 118.2. 

The ores of tin, although this metal has been known 
from very early times, occur in but few localities, and 
the metallic tin is not found in nature. The chief 
European sources of tin are the Cornish mines, where 
it is found as tin dioxide or tinstone. Tinstone is also 
met with in Malacca, and Borneo, and Mexico. In or- 
der to prepare the metal, the tinstone is crushed and 
washed, to remove mechanically the lighter portions 
of rock with which it is mixed, and the purified ore is 
then placed in a reverberatory furnace with anthracite 
or charcoal and a small quantity of lime. The oxide is 
thus reduced, and the liquid metal, together with the 
slag, consisting of silicate of lime, falls to the lower 
part of the furnace. The blocks of tin, still impure, 
are then refined by gradually melting out the pure tin, 
leaving an impure alloy behind. English tin generally 
contains traces of arsenic, copper, and other metals, 
that imported from Banca is nearly chemically pure. 

Tin possesses a white color resembling that of silver, 
it is soft, malleable, and ductile, but possesses little 
tenacity, a wire two mms. in diameter breaking with 
a weight of sixteen kilos. When bent, pure tin emits 
a peculiar crackling sound. Tin melts at 235°, and is 
not sensibly volatile. Tin does not lose its lustre on 
exposure to the air, whether dry or moist, at ordinary 



TIN 257 

temperature, but if strongly heated it takes fire, and a 
white powder of stannic oxide is formed. Hydro- 
chloric acid dissolves tin with the evolution of hydro- 
gen and the formation of stannous chloride. Nitric 
acid also attacks the metal with great energy, nitrous 
fumes being given off and stannic oxide being left as 
a white powder. 

Tin can easily be distinguished in solution by the 
formation of a splendid purple color called purple of 
cassius, formed when gold chloride is added to a dilute 
solution of stannous chloride. Tin is also easily re- 
duced before the blowpipe in the form of white mal- 
leable beads. Tin is largely used in the arts for cover- 
ing and thus protecting iron plates, or for tin-plating. 

Tin and Chlorine Compounds. 
Stannic Chloride: Formula — SnCl 4 . 

Stannic chloride is obtained by passing chlorine gas 
over metallic tin, it is a colorless liquid, boiling at 
120° Centigrade, and having a vapor density of 9.2. 
It fumes strongly in the air, and forms a crystalline 
hydrate, when a small quantity of water is added, 
which easily dissolves in an excess. Stannic chloride 
is also used by dyers, and is prepared for this purpose 
by dissolving tin in cold nitro-hydrochloric acid. 

Stannous Chloride: Formula — SnCl 2 . 

Stannous chloride is obtained by dissolving tin in 
hydrochloric acid, and separates out in needle-shaped 
crystals, when the solution is concentrated. Stannous 
chloride is termed tin salts in commerce, it is largely 



258 ELEMENTARY CHEMISTRY 

manufactured for the calico-printer and dyer, who use 
it as a mordant. 



Tin and Oxygen Compounds. 

Tin Dioxide, Sn0 2 and Tin Monoxide, SnO. 



Stannic or Tin Dioxide: Formula — SnO 



2' 



Tin dioxide occurs native as tinstone, and it can be 
prepared as a hydrate in two conditions, possessing 
totally different properties. If tin be oxidized by nitric 
acid, hydrated stannic oxide is produced as a white 
powder insoluble in acids. On the other hand if to a 
solution of stannic chloride an alkali be added, a white 
precipitate is formed of hydrated stannic oxide, which 
is readily soluble in acids. Both of these varieties of 
hydrated stannic oxide form salts, the insoluble com- 
pound having been termed metastannic, and the solu- 
ble compound stannic acid. Sodium stannate is formed 
by boiling stannic oxide with soda and is largely used 
in calico-printing as a mordant. 

Stannous or Tin Monoxide: Formula — SnO. 

This is a black powder prepared by heating the 
stannous hydrate in an atmosphere of carbonic acid, it 
rapidly absorbs oxygen from the air, passing into stan- 
nic oxide. The hydrate falls as a white powder when 
a solution of a stannous salt is added to an alkaline 
carbonate. 



TUNGSTEN 259 

Tin and Sulphur Compounds. 

Stannous Sulphide: Formula — SnS. 

Stannic Sulphide: Formula — SnS 2 . 

Of the sulphides of tin, stannous sulphide and stan- 
nic sulphide are the most important. The former is 
blackish-grey, and the latter a bright yellow crystalline 
powder known as mosaic gold is soluble in alkaline 
sulphides. 

TUNGSTEN. 
Symbol— W. Atomic Weight— 182.7. 

This metal, occurs in tolerably large quantities com- 
bined with ferrous oxide in the mineral wolfram, and 
also with lime as scheelite. The metal has only been 
obtained as a greyish-black powder. Tungsten is em- 
ployed in the arts, the addition of a small quantity im- 
parts a great degree of hardness and other valuable 
qualities to steel. 

Tungsten and Oxygen Compounds. 

Two oxides of tungsten are known, Tungsten diox- 
ide, W0 2 , and Tungsten trioxide, W0 3 . The former 
of these is obtained as a brown powder by heating the 
trioxide in an atmosphere of hydrogen, the latter, some- 
times called tungstic acid, is obtained as an insoluble 
yellow powder by heating the native calcium tungstate 
with nitric acid. Tungsten trioxide forms a variety of 
somewhat complicated salts. The sodium compound is 
soluble, and has been used to add to the starch em- 
ployed to stiffen light fabrics, the tungstate rendering 
the fabric uninflammable. 



260 ELEMENTARY CHEMISTRY 

ZINC. 

Symbol— Zn. Atomic Weight— 64.9. 

Zinc is an abundant and useful metal, closely re- 
sembling magnesium in its chemical characters, but it 
is much more easily extracted from its ores than this 
latter metal. The chief ores of zinc are the sulphide or 
blende, the carbonate or calamine, and the red oxide. 
In order to extract the metal, the powdered ore is 
roasted, or exposed to air at a high temperature, so as 
to convert the sulphide or carbonate into oxide, the 
roasted ore is then mixed with fine coal or charcoal 
and strongly heated in crucibles or retorts of peculiar 
shape, the zinc oxide is reduced by the carbon, carbon 
monoxide comes off, and the metallic zinc distills over, 
and is easily condensed. 

Zinc is a bluish-white metal, exhibiting crystalline 
structure, it is brittle at the ordinary temperature, but 
when heated to about 130°, it may be rolled out or 
hammered with ease, whilst if more strongly heated to 
200°, it is again brittle, and may be broken up in a 
mortar. Zinc melts at 423°, and at a bright red heat 
it begins to boil, and volatilizes, or if air be present it 
takes fire and burns with a luminous greenish flame, 
forming zinc oxide. Zinc is not acted upon by moist 
or dry air, and hence it is largely used in the form of 
sheets, and is employed as a protecting covering for 1 
iron, which when thus coated is said to be galvanized. 
Zinc easily dissolves in dilute acids with evolution of 
hydrogen, and it is thus used as the oxidizable portion 
of the galvanic battery. Brass is a useful alloy of one 
part of zinc and two of copper, German silver is an 
alloy of zinc, nickel, and copper. 



ZINC 261 



Zinc and Oxygen Compounds. 
Zinc Oxide: Formula — ZnO. 

Zinc oxide is the only known compound of this metal 
with oxygen, and is obtained by burning the metal, or 
by precipitating a soluble zinc salt with an alkali, and 
heating the precipitate. Zinc oxide is an insoluble 
white amorphous powder, which when heated becomes 
yellow, but loses this color on cooling, it dissolves easily 
in acids, giving rise to the zinc salts. 

The salts of zinc can be distinguished by the solu- 
bility of the oxide in excess of both potash and am- 
monia, by the white sulphide insoluble in acetic acid, 
and by the green color which a solution of cobalt chlor- 
ide imparts to zinc salts when heated in the blowpipe. 

Zinc, Carbon and Oxygen Compounds. 
Carbonate of Zinc: Formula — ZnCO r 

Zinc carbonate is an insoluble substance, occurring 
native as calamine, it cannot be prepared by precipi- 
tating a solution of zinc salt by an alkaline carbonate, 
as a quantity of oxide is precipitated along with the 
carbonate. 

Chloride of Zinc: Formula — ZnCl 2 . 

Zinc chloride is a white soluble deliquescent sub- 
stance, formed by burning zinc in chlorine, or by dis- 
solving the metal in hvdrochloric acid. 



262 ELEMENTARY CHEMISTRY 

Sulphate of Zinc: Formula — ZnS0 4 . 

Zinc sulphate is a soluble salt, crystalizing in long 
prisms, and commonly called white vitriol. This salt is 
isomorphous with magnesium sulphate, and, like the 
latter salt, it forms a series of double salts with al- 
kaline sulphates. 

Zinc Sulphide: Formula — ZnS. 

Zinc sulphide occurs as a crystalline mineral called 
blende, generally colored, from presence of iron and 
other impurities. It is obtained artificially as a white 
gelatinous precipitate, insoluble in acetic, but soluble 
in a mineral acid, formed when an alkaline sulphide is 
added to a soluble zinc salt. 



DETERMINATION BY WEIGHT. 

As it is the aim of the chemist to examine the prop- 
erties of the elements and their compounds, and as the 
weight-determination of a substance is of the greatest 
importance, it becomes necessary for him to ascertain 
with great precision the proportion by weight in which 
these several elements combine, as well as that in which 
any one of them occurs in a given compound, and for 
this purpose the Balance is employed. By means of 
this instrument the weight of a given substance is com- 
pared with the unit of weight. It consists essentially 
of a light but rigid brass beam (Fig. 40), suspended 
on a fixed horizontal axis situated at its centre ; and 
this beam is so hung as to assume a horizontal position 
when unloaded. At each end of the beam scale-pans 
are hung, one to receive the body to be weighed and 
the other for the weights. When each pan is equally 
weighted the beam must still retain its horizontal posi- 
tion or oscillate about this position, but when one pan 
is more heavily weighted than the other, the beam will 
incline on the side of the heavier pan. The balance is, 
therefore, a lever with equal arms, and it is evident that 
the weight of the substance relative to the unit weight 
employed is the sum of the weights necessary to bring 
the balance into equilibrium. The two important requi- 
sites in a balance are accuracy and sensibility, and 
these can only be gained by careful construction. It 
needs but little consideration to see that in a delicate 
balance the friction of the various parts must be re- 

263 



264 



ELEMENTARY CHEMISTRY 



duced to a minimum. This is usually accomplished by 
suspending the beam by means of an agate knife-edge, 
working on an agate plane, whilst the pans are at- 
tached to each end of the beam by a somewhat similar 
arrangement. The position of the axis of suspension 
relatively to the centre of gravity of the beam is like- 
wise a matter of consequence. If the axis of suspension 




Fig. 40. 



and the centre of gravity in a balance were coincident, 
the beam would remain stationary in all positions in 
which it might be placed. If the axis of suspension 
be placed below the centre of gravity the beam would 
be in a condition of unstable equilibrium. Hence the 
only case in which the balance can be used is that in 
which the point or axis of suspension is above the 
centre of gravity, for in this case alone will the beam 



DETERMINATION BY WEIGHT 265 

return to a horizontal position after making an oscil- 
lation, and in this case the balance may be considered 
as a pendulum, the whole weight of the beam and pans 
being regarded as concentrated at the centre of gravity. 
In order that the weight of the substance and the sum 
of the measuring weights in the scale-pan may be 
equal, it is evident that the axis of suspension must be 
exactly in the centre of the beam, or in other words, 
that the balance must have arms of equal length. It is 
also necessary that the balance should have great sen- 
sibility; that is, that it may be moved by the smallest 
possible weight; for this end it is likewise requisite 
that the vertical distance of the centre of gravity be- 
low the axis of suspension should be as small as pos- 
sible. As the whole weight of the instrument may be 
regarded as concentrated at the center of gravity, it 
evidently requires a less force to act at the end of the 
beam to move the instrument when the distance of the 
centre of gravity from the point of suspension of the 
balance is small, than when that distance is greater, 
inasmuch as in the latter ease the weight has to be 
lifted through a longer arc. The sensibilty of the bal- 
ance is also increased, both by increasing the length of 
the beam and by diminishing the weight of the beam 
and of the load. When, however, the beam is made 
either too long or too light it cases to be rigid, and a 
serious source of error is introduced. 

In all weighings with delicate balances it is neces- 
sary to have recourse to the method of weighing by 
vibration by which the excursions of the moving beam 
are accurately observed instead of its approach to the 
horizontal position. 

A good chemical balance, such as is commonly used 



266 ELEMENTARY CHEMISTRY 

for analytical work will indicate 0.0001 gram, when 
loaded with 50 to 100 grams in each pan. With a spe- 
cially constructed balance by combining this method of 
vibrations with that of double weighing which consists 
in reversing the position of the loads in the two pans, 
it is possible with a load of 1 kilogram in each pan to 
ascertain definitely a difference of weight of 0.0001 
gram or the 1/10,000,000 of the weight in either pan. 



DETERMINATION BY WEIGHT 



267 



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INDEX 



PAOB 

Atomic Theory 5 

Chemical Combination 15 

Combining Volumes of Gases 26 

Crystallization 29 

Determination of the Molecular Weight of Gases .... 30 

Flame and the Combustion of Hydrocarbons * 32 

Liquids 39 

Matter 41 

List of Elements 47 

Properties of Gases 48 

Properties of Solutions 50 

Volume of Gases 52 

Diffusion of Gases 53 

THE ELEMENTS. 

Aluminum 55 

Aluminum and Oxygen 56 

Aluminum and Chlorine 56 

Aluminum, Sulphur and Oxygen 56 

Antimony 57 

Antimony and Oxygen 58 

Arsenic 59 

Arsenic and Oxygen 61 

Arsenic and Sulphur. 62 

Barium 65 

Barium and Chlorine 65 

Barium and Sulphur 65 

Barium and Oxygen 66 

Bismuth . 66 

Bismuth and Oxygen 67 

Bismuth and Nitrogen 67 

Boron 68 

269 



INDEX 



Boron, Amorphous 68 

Boron, Crystalline . . , 70 

Boron and Carbon 71 

Bromine 72 

Bromine and Hydrogen 75 

Cadmium 76 

Calcium 76 

Calcium and Chlorine 76 

Calcium and Oxygen 77 

Calcium, Carbon and Oxygen 78 

Calcium, Chlorine and Oxygen 78 

Calcium and Sulphur 79 

Carbon 79 

The Diamond 81 

Graphite 81 

Charcoal 82 

Carbon and Hydrogen 83 

Carbon, Hydrogen and Oxygen . 87 

Carbon and Nitrogen 88 

Carbon, Hydrogen and Chlorine 90 

Carbon, Hydrogen and Oxygen 90 

Carbon and Oxygen 92 

Carbon and Sulphur 95 

Chlorine 98 

Chromium 105 

Chromium and Chlorine 105 

Chromium and Oxygen 106 

Cobalt 107 

Cobalt and Chlorine 108 

Cobalt and Oxygen 108 

Copper 109 

Copper and Chlorine Ill 

Copper and Oxygen Ill 

Copper, Sulphur and Oxygen 112 

Copper and Sulphur 112 

Fluorine- 113 

Gold .: 114 

Gold and Oxygen 115 

Hydrogen '. 115 

Hydrogen and Chlorine 120 



INDEX 271 

Hydrogen and Fluorine 123 

Hydrogen and Oxygen 125 

Hydrogen and Sulphur 146 

Hydrogen, Carbon and Nitrogen 149 

Hydrogen, Boron and Oxygen 151 

Hydrogen, Bromine and Oxygen 152 

Hydrogen, Nitrogen and Oxygen 154 

Hydrogen, Sulphur and Oxygen 161 

Iron 167 

Iron and Oxygen 169 

Iron Carbon and Oxygen 173 

Iron and Chlorine 173 

Iron, Sulphur and Oxygen . . . .* 174 

Iron and Sulphur 174 

Iodine 174 

Lead 178 

Lead and Chlorine 180 

Lead, Chromium and Oxygen 180 

Lead and Iodine 180 

Lead and Oxygen 181 

Lead, Sulphur and Oxygen 182 

Lithium 182 

Magnesium 183 

Magnesium, Carbon and Oxygen 183 

Magnesium and Chlorine 184 

Magnesium and Oxygen 184 

Magnesium, Sulphur and Oxygen 184 

Manganese 185 

Manganese and Oxygen 185 

Manganese, Potassium and Oxygen 186 

Mercury 185 

Mercury and Chlorine 188 

Mercury and Oxygen 189 

Mercury and Sulphur . 189 

Molybdenum 189 

Nickel 190 

Nickel and Oxygen 190 

Nitrogen 190 

Nitrogen and Chlorine 195 

Nitrogen and Hydrogen 196 



272 INDEX 



Nitrogen and Oxygen KA 202 

Oxygen 214 

Ozone 218 

Platinum V 220 

Platinum and Oxygen 221 

Platinum and Chlorine 221 

Potassium .222 

Potassium and Chlorine ._ 224 

Potassium and Iodine 224 

Potassium and Oxygen 224 

Potassium, Oxygen and Hydrogen 225 

Potassium, Carbon and Oxygen 225 

Potassium, Chlorine and Oxygen 226 

Potassium, Hydrogen, Carbon and Oxygen 226 

Potassium, Nitrogen and Hydrogen 227 

Potassium, Sulphur and Oxygen 227 

Phosphorus 228 

Silicon .233 

Silicon and Carbon 235 

Silicon and Oxygen 236 

Silicon, Oxygen and Hydrogen 237 

Silver .239 

Silver and Chlorine 241 

Silver and Oxygen 241 

Silver, Nitrogen and Oxygen 242 

Sodium 242 

Sodium, Carbon and Oxygen 243 

Sodium and Chlorine 244 

Sodium and Oxygen 244 

Sodium, Hydrogen and Oxygen 245 

Sodium, Hydrogen, Carbon and Oxygen 245 

Sodium, Nitrogen and Oxygen 246 

Sodium, Sulphur and Oxygen 246 

Sulphur .246 

Sulphur and Oxygen 251 

Tin 256 

Tin and Chlorine 257 

Tin and Oxygen : 258 

Tin and Sulphur 259 

Tungsten 259 



. 



INDEX 273 

Tungsten and Oxygen 259 

Zinc' 260 

Zinc and Oxygen 261 

Zinc, Carbon and Oxygen 261 

Zinc, Sulphur and Oxygen 262 

Zinc and Sulphur ..262 

Determination by weight 263 

Metric System of Weights and Measures 267 

COMPOUNDS. 

Alumina 56 

Aluminum Chloride 56 

Aluminum Sulphate . 56 

Antimony Trioxide 58 

Antimony Pentoxide 58 

Arsenious Oxide 61 

Arsenic Disulphide 62 

Arsenic Subsulphide 62 

Barium Chloride 65 

Barium Sulphate 65 

Barium Monoxide 66 

Barium Dioxide 66 

Bismuth Trioxide 67 

Bismuth Pentoxide 67 

Bismuth Sulphide 67 

Bismuth Nitrate 67 

Bismuth Trichloride . 67 

Boron Carbide 71 

Hydrobromic Acid 75 

Calcium Chloride 76 

Calcium Oxide or Lime 77 

Calcium Carbonate 78 

Chloride of Lime 78 

Calcium Sulphate . , 79 

Methane or Marsh Gas 84 

Acetylene 85 

Ethylene or Olefiant Gas 86 

Ether or Diethyl Ether 87 

Cvanogen Gas 88 



274 INDEX 

Chloroform 90 

Methyl or Wood Alcohol 90 

Ethyl or Absolute Alcohol 91 

Carbon Dioxide or Carbonic Acid Gas. 92 

Carbon Monoxide or Carbonic Oxide Gas 94 

Carbon Bisulphide 95 

Chromic Chloride 105 

Chromium Monoxide 106 

Chromium Sesquioxide , 106 

Chromium Trioxide 106 

Chromic Acid and the Chromates 107 

Cobalt Chloride 108 

Cobalt Monoxide 108 

Cobalt Sesquioxide 108 

Cobalt Oxide .108 

Chloride of Copper Ill 

Cuprous or Red Oxide of Copper Ill 

Black Oxide of Copper 112 

Sulphate of Copper 112 

Sulphide of Copper 112 

Gold Suboxide 115 

Gold Trioxide 115 

Hydrochloric or Muriatic Acid 120 

Hydrofluoric Acid 123 

AVater or Hydrogen Monoxide 125 

Peroxide of Hydrogen or Hydrogen Dioxide 125 

Hydrogen Sulphide or Sulphuretted Hydrogen .... 146 

Hydrogen Disulphide 149 

Hydrocyanic or Prussic Acid 149 

Boric or Boracic Acid 151 

Hypobromus Acid 152 

Bromic Acid 153 

Nitric Acid or Hydrogen Nitrate 154 

Nitrous Acid 159 

Sulphuric Acid or Oil of Vitriol 161 

Sulphurous Acid 165 

Iron Monoxide 169 

Iron Sesquioxide or Ferric Oxide 169 

Black Oxide of Iron 170 

Carbonate of Iron 173 



INDEX 275 

Ferrous Chloride 173 

Sulphate of Iron 171 

Ferrous Sulphide 174 

Lead Chloride 180 

Lead Chromate 180 

Lead Iodide . . . 180 

Lead Monoxide 181 

Lead Dioxide 181 

Lead Oxide . 181 

Lead Sulphate 182 

Magnesium Carbonate 183 

Magnesium Chloride 184 

Magnesium Oxide 184 

Magnesium Sulphate 184 

Manganese Monoxide 185 

Manganese Sesquioxide 186 

Manganese Dioxide , 186 

Manganic Acid 186 

Permanganate of Potash 186 

Mercuric Chloride 188 

Mercurous Chloride of Calomel 188 

Mercuric Oxide 189 

Mercuric Sulphide or Cinnabar 189 

Nickel Monoxide 190 

Nickel Sesquioxide 190 

Nitrogen Chloride 195 

Ammonia 197 

Hydrazine or Diamide 200 

Azoimide or Hydrazoie Acid 201 

Nitrogen Monoxide 208 

Nitrogen Dioxide 208 

Nitrogen Trioxide : 208 

Nitrogen Peroxide 208 

Nitric Anhydride 208 

Platinum Dichloride 221 

Platinum Tetrachloride 221 

Potassium Chloride 224 

Potassium Iodide 224 

Potassium Monoxide 224 

Potassium Dioxide .224 



276 INDEX 

Potassium Tetroxide 224 

Potassium Hydroxide 225 

Potassium Carbonate 225 

Chlorate of Potash 226 

Bicarbonate of Potash 226 

Potassium Nitrate or Saltpeter 227 

Sulphate of Potash : 227 

Silicon Carbide 235 

Silicon Dioxide 236 

Silicic Acid 237 

Silver Chloride 241 

Silver Suboxide . . 241 

Silver Monoxide '. 241 

Silver Dioxide 242 

Nitrate of Silver .242 

Sodium Carbonate or Soda-ash 243 

Chloride of Sodium 244 

Sodium Oxide 244 

Sodium Dioxide 244 

Sodium Hydroxide 245 

Sodium Bicarbonate 245 

Sodium Nitrate 246 

Sulphate of Soda .246 

Sulphur Dioxide 252 

Sulphur Trioxide ....'. 254 

Stannic Chloride 257 

Stannous Chloride* 257 

Stannic or Tin Dioxide 258 

Stannous of Tin Monoxide 258 

Stannous Sulphide 259 

Stannic Sulphide 259 

Tungsten Dioxide 259 

Tungsten Trioxide 259 

Zinc Oxide 261 

Zinc Carbonate 261 

Zinc Chloride 261 

Zinc Sulphate 262 

Zinc Sulphide 262 



THE MOST IMPORTANT BOOK ON ELECTRICAL CONSTRUCTION 

WORK FOR ELECTRICAL WORKERS EVER PUBLISHED. 

NEW 1904 EDITION. 

MODERN WIRING 

DIAGRAMS AND DESCRIPTIONS 

A Hand Book of practical diagrams and 
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By HENRY C. HORSTMANN and 
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This grand little volume not only tells 
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Partial Table of Contents 
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MODERN ELECTRICAL 
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TTThIS book treats almost entirely of practical electrical 
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ELECTRICAL 
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Easy Electrical Experiments 
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MB— — — I 

Or Oil and Water Color 
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DYNAMO TENDING 




ENGINEERS 

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VICTOR H. TOUSLE Y, 
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